Up until now we have been discussing only the elemental forms of atoms which are neutrally charged. This is because the number of electrons (negative in charge) is equal to the number of protons (positive in charge).

This one-to-one ratio of charges is not, however, the most common state for many elements. Deviations from this ratio result in charged particles called ions.

Lower energy configurations are more stable, so things are naturally drawn toward them. For atoms, these lower energy states are represented by the noble gas elements.

This makes them stable and unreactive. They are already at a low energy state, so they tend to stay as they are.

This instability drives them toward the lower energy states represented by the noble gases that are nearby in the periodic table. In these lower energy states, the outermost energy level has eight electrons (an “octet”).

There are two ways for an atom that does not have an octet of valence electrons to obtain an octet in its outer shell. One way is the transfer of electrons between two atoms until both atoms have octets.

Those that lose electrons become positively charged, and those that gain electrons become negatively charged. Recall that atoms carrying positive or negative charges are called ions.

If an atom has lost one or more electrons, it is positively charged and is called a cation. Because opposite charges attract (while like charges repel), these oppositely charged ions attract each other, forming ionic bonds.

The second way for an atom to obtain an octet of electrons is by sharing electrons with another atom. These shared electrons simultaneously occupy the outermost shell of both atoms.

Covalent bonding and covalent compounds will be discussed in Chapter 4 “Covalent Bonding and Simple Molecular Compounds”. For each element on the periodic table, it is possible to predict the number of valence shell electrons that they will contain.

The main group elements are numbered IA to VIIIA, and their number of valance shell electrons correspond to their group number (Fig 4.1). For example all of the elements in the halogen family belong to group VIIA and correspondingly have 7 electrons in their valence shell.

Figure 4.1 Periodic Table of the Elements. Main group, transition and inner transition elements are indicated.

The graphical notation used for valence electrons is called an Electron-Dot Symbol. To draw an electron dot symbol, start with the abbreviation for the element of interest as the center, signifying the nucleus of the atom.

Students often want to place these electron dots around the element randomly, but it is useful to use the four cardinal directions as a guide. First place single electrons around the atom in each of the four cardinal positions until you run out of electrons.

Note that the noble gases have complete octets and will have a total of 8 electrons in their valence shell (Fig 4.2).

Figure 4.2 Electron Dot Symbols. Electron Dot Symbols are shown for Carbon through Neon on the Periodic Table.

Overall, the periodic table can be used as a guide for determining the number of valence electrons for each element (Fig 4.3).

Figure 4.3 Periodic Table with Electron Dot Symbols. Electron dot symbols are drawn above each family or group of elements on the periodic table, where X indicates any element within that family or group.

The elements on the right side of the periodic table, nonmetals, gain the electrons necessary to reach the stable electron configuration of the nearest noble gas. Elements on the left side of the periodic table, metals, lose the electrons necessary to reach the electron configuration of the nearest noble gas.

Group IA elements form ions with a +1 charge. They lose one electron upon ionization, moving into the electron configuration of the previous noble gas.

The sodium ion has one fewer electron than it has protons, so it has a single positive charge and is called a cation.

Sodium tends to lose it’s valence shell electron in the third shell during ionic bond formation. It is left with a full octet in the second shell and now has the electron configuration of neon.

However, there are now only 10 electrons within the electron cloud, resulting in a net positive (+1) charge. Upon losing that electron, the sodium ion now has an octet of electrons from the second principal energy level.

The term isoelectronic refers to an atom and an ion of a different atom (or two different ions) that have the same electron configuration. The sodium ion is isoelectronic with the neon atom.

Overall, Group IA elements will lose one electron to reach the electron state of the noble gas preceding them in the periodic table. Note that the nucleus of the atom remains unchanged and thus, the identity of the ion is also unchanged.

Thus, it retains its identity as the element, sodium even when it has undergone the loss of an electron. Similarly, Group IIA elements lose two valence electrons to form ions with a +2 charge and Group IIIA elements lose three electrons to form ions with a +3 charge.

While hydrogen is in the first column, it is not considered to be an alkali metal, and so it does not fall under the same classification as the elements below it in the periodic table. This is because hydrogen is very small and can only house a total of 2 electrons to become filled.

Thus, instead of following the octet rule, it reaches greater stability by gaining a “duet” of electrons through bonding with other atoms. Thus, hydrogen can form both covalent bonds and ionic bonds, depending on the element that it is interacting with.

Note, that hydrogen only has one electron to begin with, so when it loses an electron in the ionized state, there is only a single proton left in the nucleus of the atom. Thus, when hydrogen is ionized to H+ it is often referred to as a proton.

In this case, the H– anion is named using standard convention forming the hydride ion. During the ionization of hydrogen, the H+ state is more common than the H– state.

Acids are defined as compounds that donate H+ ions in aqueous solutions. Cations are named very simply by following the element name with the word ‘ion’.

Elements on the other side of the periodic table, the nonmetals, tend to gain electrons in order to reach the stable electron configurations of the noble gases that come after them in the periodic table. Group VIIA elements gain one electron when ionized, obtaining a -1 charge.

This gives it a single negative charge, and it is now a chloride ion (Cl–). note the slight change in the suffix (-ide instead of -ine) to create the name of this anion.

Fig 4.5 The Formation of a Chloride Ion. On the left, a chlorine atom has 17 electrons.

“An ion is a small electrically charged particle. Ions are single charged atoms (simple ions) or small charged “molecules” (polyatomic ions).”.

Ions are formed as a result of this addition or removal of electrons from the atom. Negative ions, known as anions, form when an atom gains electrons and now has more electrons than protons, indicating that the number of protons and electrons is not balanced.

Atoms do this so that they can have a full outer shell of electrons and be more energy stable.

It can be used to predict whether or not an atom can bond with another atom. An atom’s charge is related to its valence electrons or oxidation state.

There are only two types of ions- Cations and anions. Cations are positively charged particles or ions and anions are negatively charged particles or ions.

The group number minus eight equals the number of charges on an ion formed by a nonmetal. When writing an ion symbol, the one- or two-letter element symbol comes first, followed by a superscript.

Because neutral atoms have a charge of zero, no superscript is required. Since the Group 8A elements have a full octet of eight valence electrons in their highest-energy orbitals, they have a low tendency to gain or lose electrons to form ions or share electrons with other elements in covalent bonds.

Therefore, ions do not usually exist on their own but will bind with ions of opposite charges to form a crystal lattice. The resulting compound is called an ionic compound, and is said to be held together by ionic bonding.

In this tutorial, you will be introduced to ionic radius trends on the periodic table of elements. You will also be introduced to the concepts that contribute to ionic radius, including how to find it.

When an atom on the periodic table loses an electron it becomes a cation. In contrast, when an atom in the periodic table gains an electron it becomes an anion.

For example, the metals in group 1A (Alkali Metals) all have a +1 charge meaning they want to give away an electron. The non-metals tend to become anions.

So in the picture below the ionic radius would be found by measuring the distance between the yellow and blue protons in the nucleus. To the red electrons at the outermost orbital.

The ionic radius of an ion is measured when the atom is in a crystal lattice structure. The ionic radius is half of the distance between two ions that is just touching one another.

So the ions of an atom are treated as if they were spheres. The ionic radius can easily be a little smaller or larger than the atomic radius, which is the radius a neutral atom of the element possesses.

Moreover, the ionic radius is tricky to measure since it depends on the varying factors of the environment in which the ion is located. It depends on the coordination number, or the number of atoms, ions, or molecules that a central atom or ion holds as its nearest neighbors in a complex or coordination compound.

The ionic radius is generally calculated by estimating the distance between the two nuclei and dividing it according to the atomic sizes. Ionic radius is generally measured in picometer (pm) or nanometer (nm).

Vice Versa for anions since anions gain electrons to the outermost shell they usually will have bigger radii than their parent neutral atom. Ex.

The atomic radius of the chlorine atom (Cl) is 79 pm, and the ionic radius of the chlorine ion (Cl-1) is 167 pm. As you move down the periodic table additional electrons are being added, which in turn causes the ionic radius to increase.

Congruent with the trend of increasing atomic radius as you move down the periodic table. As you move across the periodic table more electrons, protons, and neutrons are being added.

This is because as you move over a row on the periodic table the ionic radius decreases for metals forming cations. The ionic radius increases for nonmetals as the effective nuclear charge decreases.

Notice as you move to the right on the periodic table the ionic radius size is decreasing.

1 UNIT: Nomenclature Objectives: Lesson 1 of 3 You will learn which groups on the periodic table lose or gain electrons to become cations and anions You will learn what an ionic compound is and how they form You will learn how to write formulas for Ionic Compounds You will learn how to write the names of Ionic Compounds TOPIC: Ionic Compounds.

) +1 + 1 = 0 What kind of charge will balance out a charge of -1. Consider the compound NaCl, How many Na atoms are present.

3 Group Determines Charge Remember, groups of elements are in vertical columns and have similar chemical properties One of the reasons for this, is that the elements within the group form the same type of Ion 1 6 2 109453 13 712811 18 14161517 For example, group 1 alkali metals like to lose a single electron and therefore always form +1 Cations Group 2 alkali earth metals always lose two electrons and therefore always from +2 Cations Group 13 elements like Aluminum lose three electrons and always form +3 Cations Group 16 elements like to gain two electrons and therefore always from -2 Anions Group 17 elements, the halogens, like to gain one electron and therefore always from -1 Anions.

5 What determines the charge of an ion.

7 – – – – – – – – – – – – – – – – Chlorine Atom Sodium Atom Bond +1 Ionic Compounds form between a cation and an anion Ionic Compounds form when a metal atom loses an electron(s) and transfers it to a nonmetal atom For example, consider the metal Sodium (Na) and the non-metal chlorine (Cl) When these two atoms bond to make a compound, the sodium loses one electron and transfers it to chlorine +1 Salt Grain or Cube Ionic Compounds Notice after losing an electron, sodium becomes a cation and chlorine becomes anion The result is a neutral Ionic Compound where the ionic charges are balanced out or neutral Table Salt or sodium chloride (NaCl) is a perfect example of an ionic compound Ionic Compound forms between a positive cations sand a negative anions.

A neutrally charged compound that contains a cation (positive ion) and anion (negative ion) Example: Sodium Chloride: Na +1 + Cl -1  NaCl Iron (III) oxide: Fe +3 + O -2  Fe 2 O 3.

The lowest ratio of atoms in an ionic compound Incorrect: Na 2 Cl 2 Correct: NaCl. 11 Step #CationAnion 1.

Number of atoms Overall Charge 3. Formula Na +1 Cl -1 1.

Add atoms/ions until the overall charge is neutral or zero Notice, the +1 charge of sodium is balanced by the -1 charge of chlorine creating an overall charge of zero 3.

12 Step #CationAnion 1. Charge 2.

Formula Ca +2 Cl -1 1. Determine the charge for the cation and anion For example, the Calcium atom/cation has a +2 charge and Chlorine atom/anion has a -1 charge 2.

Once it is neutral, write the ionic formula, cation first and the anion second, use subscripts to show how many atoms make up the compound Now that it is neutral (+2 -1 -1 = 0) we can write the formula CaCl 2 as CaCl 2 which tells us the compound is made up of one calcium atom and one chlorine atom Write the formula for the ionic compound Calcium Chloride.

13 Step #CationAnion 1. Charge 2.

Formula Fe +3 O -2 1. Determine the charge for the cation and anion The roman numeral tells us the Iron atom/cation has a +3 charge and the Oxygen atom/anion has a -2 charge 2.

Once it is neutral, write the ionic formula, cation first and the anion second, use subscripts to show how many atoms make up the compound Now that it is neutral (+3+3 -2 -2-2 = 0), we can write the Fe 2 O 3 formula as Fe 2 O 3 which tells us the compound is made up of two Iron atoms and three Oxygen atoms Write the formula for the ionic compound Iron (III) Oxide.

14 What are the Steps for Writing Formulas for Ionic Compounds. 1.

Add atoms/ions until their overall charge is neutral or equal to zero 3. Once it is neutral, write the ionic formula with the cation first and the anion second, use subscript numbers to show how many atoms make up the compound Step #CationAnion 1.

Number of atoms Overall Charge 3. Formula Na +1 Cl -1 +1 Na Cl Cl 0 (neutral).

16 When naming ionic compounds there are certain steps you must use These steps will help you throughout this unit and the rest of the course: 1. The cation is always named first and the anion second For Example: for NaCl, we write Sodium Chloride 2.

The anion takes it’s name by taking the first part of the element name and adding -ide to the end For Example: The Cl – anion is called Chloride & the O 2- ion is called oxide 4. If the cation has more than one charge (it’s a transition metal) use roman numerals to indicate charge For Example: Iron can be either a Fe +2 or Fe +3 cation, so we write either Iron (II) or Iron(III) to indicate charge.

The cation is always named first and the anion second 2. The cation takes it name from the name of the element 3.

If the cation has more than one charge (it’s a transition metal) use roman numerals to indicate charge. 18 Practice: names Write the names for the following compounds below: 1)NaF 2)MgCl 2 3)CaS 4)Na 2 O 5)NiCl 2 6).

1 Modern Atom & Periodic Table. 2 Topic: Mass #’s, Atomic #’s, & IsotopesUnit: Modern Atom & Periodic Table Topic: Mass #’s, Atomic #’s, & Isotopes Objectives: Day 1 of 4 To understand the basic components of an atom and how we arrived at this conclusion To understand that an atom is mostly empty space To understand the difference between the mass number and the atomic number To understand how an isotope is different that of a neutral atom.

What do you think atoms look like.

What do you think happens to atoms during a chemical reaction.

4 Modern Atom e- e- The modern atom has: Electron Proton NeutronNegatively charged particle that occupies the space around the nucleus (mass = 5.48 x10-5amu) Proton Positively Charged particle in a Nucleus (mass = 1 amu) Neutron A neutral particle in a Nucleus with the same mass as a proton (mass = 1 amu) e- Nucleus size is not scale.

5 Actual Size of the NucleusIf an atom was the size of baseball stadium, the nucleus would the size of a pea on the pitcher’s mound, and the electron “cloud” would occupy the size of the stadium. 6 What is the modern atom made of.

7 Modern Atom We now know that the modern atom is made up of Protons, Neutrons and Electrons A proton and neutron has the about the same mass of 1 amu (1.66 x gms) An electron which occupies the empty space around the nucleus has a mass of x10-5amu.

How many neutrons Every element has a Mass Number and an Atomic number The Atomic Number is the number of protons The mass number is the sum of protons and neutrons Let’s look at the element Helium The atomic number (#) is the number of protons. 10 What is the difference between the Atomic Number and the Mass Number.

11 Helium Atom e- e- The HELIUM atom has: 2 protons & 2 neutronsAtomic # = 2 Mass # = 4 What is the charge in the nucleus is Helium has 2 protons. How many electrons are needed to balance a charge of +2 e- P+ N N P+ e-.

13 One Isotope of HELIUM has: But, if we add two neutrons,Isotopes This Isotope is called Helium-6 One Isotope of HELIUM has: 2 protons & 4 neutrons Atomic # = 2 Mass # = 6 Remember, neutral helium atom has 2 protons and 2 neutrons This is written as: He This is written as: 6 2 But, if we add two neutrons, an isotope forms.

14 Isotopes of Hydrogen H 3 2 1 1 N e- Tritium Deuterium P+ N. 15 Protons Determine the ElementCarbon-14 Carbon-12 6 protons 8 neutrons 6 protons 6 neutrons An element doesn’t change just because it has more neutrons, Carbon is still carbon, so long as it has 6 protons Question: What weighs more.

16 U Practice: Determine the amount of Neutrons in the isotope: 238 92A – Z = # of neutrons = # of neutrons 146 = # of neutrons. 17 What are Isotopes.

18 C Practice: Determine the amount of Neutrons in the isotope: 14 6A – Z = # of neutrons = # of neutrons 8 = # of neutrons. 19 Practice: Complete the table: 207 82 Mass #:.

Atomic #:.

Pb 82 Mass #:.

Atomic #:.

Protons:.

Neutrons:.

20 Practice: Complete the table: 17 17 20 17 37 Element/Ion Atomic NumberNumber of Protons Number of Neutrons Number of Electrons Mass Number (protons + neutrons) 17 17 20 17 37. 21 Practice: Na Complete the table: Atomic Number Number of ProtonsElement Symbol Atomic Number Number of Protons Number of Neutrons Number of Electrons Mass Number Isotope.

Element Symbol Atomic Number Mass Number Number of Protons Number of Neutrons Number of Electrons Isotope. Yes or No.

22 Cl Summarize: 37 17 What is the modern atom made of. What makes up most of the mass (weight) of an atom.

How does the size of the nucleus compare to the size of the whole atom. An isotope is…… 37 17 Cl.

24 Quickwrite Answer one of the questions below 1-2 sentences:Why do you think the periodic table is organized the way it is.

How do you think the elements vary or change throughout the periodic table.

If you can, try listing some physical properties of metals. 25 Periodic Table In every chemistry room, you can expect to see a periodic table The periodic table is listed in order of increasing atomic number It is also arranged in vertical and horizontal columns.

27 Groups Elements with similar chemical Properties that lie in the same vertical column are called groups Groups are often referred to by the number over the column Two or more groups make up a family. 28 Groups Groups are often referred to by the number over the column 1718 1 2 Groups are often referred to by the number over the column 13 14 15 16 3 4 5 6 7 8 9 10 11 12.

Elements with similar chemical ________ that lie in the same vertical column properties Answer Bank Two Properties Right Electricity Separate gases. 30 Alkali Metals Group 1 Alkali metals are elements such as Lithium, Sodium, and Potassium These metals share similar chemical properties with each other They are also very reactive and almost always bonded to another element (Ex.

31 the Alkaline Earth MetalsGroup 2 Alkaline earth metals include elements such as Beryllium, Magnesium, and Calcium These metals share similar chemical properties In group 2 we have the Alkaline Earth Metals. 32 Halogens In group 17 we havethe Halogens Group 17 Halogens include elements such as Fluorine, Chlorine and Bromine Many of these elements are gases or liquids at room temperatures.

34 Transition Metals Transition metals include many of the elements we are familiar with such as Nickel, Iron, Copper, Gold and Silver Most of the elements are Metals In groups 3-12, we have the Transition Metals. 35 What are the different types of groups.

Step 1: Identify the given atom or ion.

Step 3: Determine the number of protons or electrons from the following expressions:. Protons and Electrons in Atoms: The number of protons in the nucleus of an atom is equal to its atomic number.

Protons and Electrons in ions: The number of protons in an ion is equal to the atomic number of the corresponding atom. The number of electrons depends on the charge of the ion.

Number of electrons in an anion = Atomic number of the corresponding atom + Number of Charges. Number of electrons in a cation = Atomic number of the corresponding atom – Number of Charges.

Consult a periodic table, and determine how many electrons are in the anion {eq}Cl^- {/eq}.

b) 18. c) 10.

Step 1: Identify the given atom or ion.

Step 2: Determine the atomic number of the atom or ion.

Step 3: Determine the number of protons or electrons from the following expressions:. $$\text{Number of electrons in the anion, } Cl^- = 17+1=18 $$.

Consult a periodic table, and determine how many protons are in the cation {eq}K^+ {/eq}.

b) 18. c) 10.

Step 1: Identify the given atom or ion.

Step 2: Determine the atomic number of the atom or ion.

Step 3: Determine the number of protons or electrons from the following expressions:. $$\text{Number of protons in the cation, } K^+=19 $$.

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Up until now we have been discussing only the elemental forms of atoms which are neutrally charged. This is because the number of electrons (negative in charge) is equal to the number of protons (positive in charge).

This one-to-one ratio of charges is not, however, the most common state for many elements. Deviations from this ratio result in charged particles called ions.

Lower energy configurations are more stable, so things are naturally drawn toward them. For atoms, these lower energy states are represented by the noble gas elements.

This makes them stable and unreactive. They are already at a low energy state, so they tend to stay as they are.

This instability drives them toward the lower energy states represented by the noble gases that are nearby in the periodic table. In these lower energy states, the outermost energy level has eight electrons (an “octet”).

There are two ways for an atom that does not have an octet of valence electrons to obtain an octet in its outer shell. One way is the transfer of electrons between two atoms until both atoms have octets.

Those that lose electrons become positively charged, and those that gain electrons become negatively charged. Recall that atoms carrying positive or negative charges are called ions.

If an atom has lost one or more electrons, it is positively charged and is called a cation. Because opposite charges attract (while like charges repel), these oppositely charged ions attract each other, forming ionic bonds.

The second way for an atom to obtain an octet of electrons is by sharing electrons with another atom. These shared electrons simultaneously occupy the outermost shell of both atoms.

Covalent bonding and covalent compounds will be discussed in Chapter 4 “Covalent Bonding and Simple Molecular Compounds”. At the end of chapter 2, we learned how to draw the electron dot symbols to represent the valence electrons for each of the elemental families.

Looking at Figure 3.1, observe the Noble Gas family of elements. The electron dot symbol for the Nobel Gas family clearly indicates that the valence electron shell is completely full with an octet of electrons.

Above, we noted that elements are the most stable when they can reach the octet state. However, it should also be noted that housing excessively high negative or positive charge is unfavorable.

You will note that for the IA, IIA, IIIA and transition metals groups, it is more economical to lose electrons (1-3 electrons) from their valence shells to reach the octet state, rather than to gain 5-7 electrons. Similarly main group columns VA, VIA, and VIIA tend to gain electrons (1-3) to complete their octet, rather than losing 5-7 electrons.

These atoms don’t like to gain or lose electrons, but tend to favor the sharing model of chemical bonding. The remaining sections of this chapter will focus on the formation of ions and the resulting ionic compounds.

Figure 3.1 Periodic Table with Electron Dot Symbols.

Figure 3.2 Ionization Within and Electric Field. (A) Depiction of St.

(B) In many high voltage applications plasma ionization is an unwanted side effect. Shown is a long exposure photograph of corona discharge on an insulator string of a 500 kV overhead power line.

Photograph depicted in a (A) by: Unknown Author. Photograph depicted in a (B) by: Nitromethane.

Elements on the left side of the periodic table, metals, lose the electrons necessary to reach the electron configuration of the nearest noble gas. Transition elements can vary in how they move toward lower energy configurations.

They lose one electron upon ionization, moving into the electron configuration of the previous noble gas. For example as shown in Figure 3.3, when a sodium (Na) atom is ionized, it loses one of its 11 electrons, becoming a sodium ion (Na+) with the electron configuration that looks like the previous noble gas, neon.

Figure 3.3 The Formation of a Sodium Ion. Sodium tends to lose it’s valence shell electron in the third shell during ionic bond formation.

Note that it still has the same number of protons (11) as the original sodium atom and retains the identity of sodium. However, there are now only 10 electrons within the electron cloud, resulting in a net positive (+1) charge.

The equation below illustrates this process. The electron configuration of the sodium ion is now the same as that of the noble gas neon.

The sodium ion is isoelectronic with the neon atom. Consider a similar process with magnesium and with aluminum:

The aluminum atom loses its three valence electrons. The Mg2+ ion, the Al3+ ion, the Na+ ion, and the elemental Ne atom are all isoelectronic.

Only larger atoms, such as lead and uranium, can typically carry larger charge states.

This gives them the electron configuration of the noble gas that comes before them in the periodic table. While hydrogen is in the first column, it is not considered to be an alkali metal, and so it does not fall under the same classification as the elements below it in the periodic table.

Thus, instead of following the octet rule, it reaches greater stability by gaining a “duet” of electrons through bonding with other atoms. Thus, hydrogen can form both covalent bonds and ionic bonds, depending on the element that it is interacting with.

Note, that hydrogen only has one electron to begin with, so when it loses an electron in the ionized state, there is only a single proton left in the nucleus of the atom. Thus, when hydrogen is ionized to H+ it is often referred to as a proton.

In this case, the H– anion is named using standard convention forming the hydride ion. During the ionization of hydrogen, the H+ state is more common than the H– state.

Acids are defined as compounds that donate H+ ions in aqueous solutions, and will be discussed in more detail in Chapter 9.

Group VIIA elements gain one electron when ionized, obtaining a -1 charge. For example as shown in Figure 3.4, chlorine (Cl), when ionized, gains an electron to reach the electron configuration of the noble gas that follows it in the periodic table, argon.

note the slight change in the suffix (-ide instead of -ine) to create the name of this anion.

Fig 3.4 The Formation of a Chloride Ion. On the left, a chlorine atom has 17 electrons.

Note that the chloride ion has now filled its outer shell and contains eight electrons, satisfying the octet rule. Group VIA elements gain two elec.

Step 1: Identify the proton number, also called atomic number, of the element on the periodic table of elements.

This can be seen in the name of the ion.

Proton Number: The proton number, also called the atomic number, of an element is the number of protons in one atom of that element. The proton number can be found on the periodic table of elements.

Ion: An ion is a charged atom. In a neutral atom, the number of protons and electrons are equal.

In a negatively charged ion, also called a anion, the number of protons is less than the number of electrons.

How many electrons does a {eq}Mg^{2+} {/eq} ion have.

Looking up Magnesium, Mg, on the periodic table, we see the proton number is 12.

This can be seen in the name of the ion.

Step 3: Determine the number of electrons:. Since the ion is positively charged, we need to subtract the charge number (2) from the proton number (12) to get:.

There are 10 electrons in a {eq}Mg^{2+} {/eq} ion.

Step 1: Identify the proton number, also called atomic number, of the element on the periodic table of elements.

Step 2: If it is an ion, determine the charge of the ion. This can be seen in the name of the ion.

We have a neutral atom of Palladium, not an ion.

Since the atom is neutral, the number of electrons is equal to the proton number.

How many electrons does a {eq}P^{3-} {/eq} ion have.

Looking up Phosphorus, P, on the periodic table, we see the proton number is 15.

This can be seen in the name of the ion.

Step 3: Determine the number of electrons:. Since the ion is negatively charged, we need to add the charge number to the proton number.

$$\text{ Number of Electrons} = 15 + 3 = 18 $$. There are 18 electrons in a {eq}P^{3-} {/eq} ion.

Lithium was first discovered in 1817 when Johan August Arfwedson, a Swedish chemist, was analyzing mineral ore. Cesium and rubidium were discovered in 1860 and 1861, respectively, by German chemists Robert Bunsen (who lent his name to the Bunsen burner) and Gustav Kirchhoff (who devised Kirchhoff’s laws for electrical current).

Sodium and potassium, two very common alkali metals, have unknown discovery dates because they have been used for so long. But scientists weren’t able to isolate the pure elements until the famous chemist Humphrey Davy in 1807.

Advertisement. Alkali metals react vigorously to water and air.

They get more reactive the further down on the periodic table you go too, with cesium and francium being so reactive that they can burst into flames simply by being exposed to the air. The elements also increase in atomic radius, decrease in electronegativity and decrease in melting and boiling points as you move down the periodic table.

Well, as it turns out, most of the alkali metals are found in nature as ions due to their high desire to react and lose that one valence electron. In their ionic form the metals are far less reactive.

1 18 Oct. 2010  Take out Homework: Week 6 Homework #6- 9 AND class work from Friday  Objective: SWBAT model and describe the structure of atoms and trends in ionization.

Draw an arrow pointing to the valence shell. How many valence electrons are there.

Write the symbol.

Do now, homework check II. Review from last week III.

Hand back papers and progress reports Homework: none. 3 Review from last week.

1 2 3 4 567 8 1+ 2+3+3-2-1- Cations Anions Lose e – Gain e -. 5 How many TOTAL electrons does an atom of Lithium (Li) have.

6 How many TOTAL electrons does an atom of Magnesium (Mg) have. Objective: SWBAT model atomic structure and relate valence shell electrons to trends on the periodic table.

7 How many TOTAL electrons does an atom of Chlorine (Cl) have. Objective: SWBAT model atomic structure and relate valence shell electrons to trends on the periodic table.

8 How many VALENCE electrons does an atom of Lithium (Li) have. Objective: SWBAT model atomic structure and relate valence shell electrons to trends on the periodic table.

9 How many VALENCE electrons does an atom of Magnesium (Mg) have. Objective: SWBAT model atomic structure and relate valence shell electrons to trends on the periodic table.

10 How many VALENCE electrons does an atom of Chlorine (Cl) have. Objective: SWBAT model atomic structure and relate valence shell electrons to trends on the periodic table.

11 Since Chlorine has 7 valence electrons, to satisfy the octet rule, it must: Objective: SWBAT model atomic structure and relate valence shell electrons to trends on the periodic table. 1.

Gain one electron 3. Lose one electron 4.

12 Since Lithium has 1 valence electron, to satisfy the octet rule, it must: Objective: SWBAT model atomic structure and relate valence shell electrons to trends on the periodic table. 1.

Gain one electron 3. Lose one electron 4.

13 Since Magnesium has 2 valence electrons, to satisfy the octet rule it must: Objective: SWBAT model atomic structure and relate valence shell electrons to trends on the periodic table. 1.

Gain 2 electrons 3. Lose 2 electrons 4.

14 Since neon (Ne) has 8 valence electrons, to satisfy the octet rule it must: Objective: SWBAT model atomic structure and relate valence shell electrons to trends on the periodic table. 1.

Gain 2 electrons 3. Lose 8 electrons 4.

15 Ions formed by group – Add to your PTE Objective: SWBAT model atomic structure and relate valence shell electrons to trends on the periodic table. 1 2 3 4 567 8 1+ 2+3+3-2-1- Cations Anions Lose e – Gain e -.

Objective: SWBAT model atomic structure and relate valence shell electrons to trends on the periodic table. 1.

Li – 3. Li 2+ 4.

17 What is the correct symbol for a Magnesium ion. Objective: SWBAT model atomic structure and relate valence shell electrons to trends on the periodic table.

Mg + 2. Mg – 3.

Mg 2-. 18 Which of the following pairings is correct.

Magnesium/Cation 2. Chlorine/Anion 3.

Sulfur/Anion 5. All of the above.

Objective: SWBAT model atomic structure and relate valence shell electrons to trends on the periodic table. 1.

Fluorate 3. Fluorite 4.

20 What is the name of the ion formed from the element Lithium. Objective: SWBAT model atomic structure and relate valence shell electrons to trends on the periodic table.

Lithium 2. Lithiate 3.

Lithide. 21 C: 19 Oct.

 Draw a picture of one atom of sulfur. Draw an arrow showing the valence shell.

Total number of electrons: ___________ 2. Number of valence electrons: ___________ 3.

________ 4. How many will it gain or lose.

Symbol with charge: _________. 22 Agenda I.

Ionic bonding notes III. Naming compounds practice IV.

Writing formulas practice VI. Ionic bonding “Speed Dating” Homework: Finish Week 6 Homework SWBAT model ionic bonding and write names and formulas for ionic compounds.

2010 Ionic Bonding  Copy in notebook: Objective: SWBAT model ionic bonding and write names and formulas for ionic compounds  Do now: On handout.

Draw an arrow showing the valence shell. Answer the questions to the right.

Total number of electrons: ___________ 2. Number of valence electrons: ___________ 3.

________ 4. How many will it gain or lose.

Symbol with charge: _________. 24 Agenda I.

Ionic bonding notes III. Naming compounds practice IV.

Writing formulas practice Homework: Finish Week 6 Homework SWBAT model ionic bonding and write names and formulas for ionic compounds.

27 Ionic Bonding Demo with marbles SWBAT model ionic bonding and write names and formulas for ionic compounds. 28 Sodium metal reacts with chlorine gas to form solid sodium chloride Na + + Cl –  NaCl Na.

30 A: 20 Oct. 2010  Take Out Homework: Week 6 Homework  Objective: SWBAT model ionic bonding and write names and formulas for ionic compounds.

magnesium chloride 2. CaO 3.

AlBr 3 5. aluminum iodide.

Do now II. Homework check III.

Ionic bonding “Speed Dating” Homework: Revise your Week 6 Homework. 32 Naming ionic compounds  Write the cation name  Name of the element  Write the anion name  Beginning of element name + -ide  Ex 1: LiCl  Ex 2: Al 2 O 3 SWBAT model ionic bonding and write names and formulas for ionic compounds.

CaF 2 2. BeO 3.

KBr 5. Rb 2 S 6.

SrCl 2 8. MgO SWBAT model ionic bonding and write names and formulas for ionic compounds.

Write each ion symbol with the charge 2. “Cross” the charges down to the bottom of the opposite element 3.

Do not write + or – Example: Ex 1: lithium chloride Ex 2: aluminum oxide Ex 3: lithium sulfide. 35 Writing formulas from names 1.

magnesium bromide 3. rubidium sulfide 4.

sodium oxide 6. magnesium sulfide 7.

aluminum iodide. 36 IONS SPEED DATING Objective: SWBAT model atomic structure and relate valence shell electrons to trends on the periodic table.

Ionic radii, what is it. Let’s figure out what we mean by ionic radii and their variation in groups and periods of the periodic table (modern eriodic table).

Ions are formed when an atom loses or gains electrons. When an atom loses an electron it forms a cation and when it gains an electron it becomes an anion.

The atomic size of a cation will be smaller than that of the parent atom. An anion is relatively larger in size than its parent atom.

For example,. Atoms and ions which consist of an equal number of electrons are considered isoelectronic species.

The radius of a cation will be smaller than that of the anion as a cation will have a greater positive charge (i.e. number of protons) so it will attract the electrons in the outermost orbital with greater force and hence the smaller size.

Let us understand the trends in the ionic radius of elements across a period with an example. In period 3 we find that the atomic radius first decreases and then suddenly increases and then again it slowly decreases.

Ions are formed as a result of the gain or loss of electrons. The formation of ions plays a vital role in any chemical reaction to form a new substance.

To know more about elements of the periodic table and its properties visit BYJU’S.

KEY CONCEPTS. All matter is made of atoms of individual elements.

Earth, however, is dominated by other elements, and 8 to 10 elements account for most of Earth’s mass. Although Earth’s crust and mantle contain the same major elements, the proportions are not the same (see histograms below).

Most of the minerals and other geological materials we see derive from the crust, but some come from the uppermost part of the mantle. In both places, oxygen and silicon are the dominant elements (Figures 2.2 and 2.3).

Other quite abundant elements in both the crust and mantle include aluminum, iron, calcium, sodium, potassium, and magnesium. However, the mantle contains much more magnesium and iron, and less silicon, than the crust.

And other generally rare elements are sometimes concentrated by geological process to make exotic minerals. The compositions of Earth’s outer layers vary somewhat laterally and vertically.

And, the composition of the shallow crust is somewhat different from the deep crust. Nonetheless, in most settings, we can expect common minerals to be made of the elements shown in the histograms above.

Oxygen and silicon are perhaps the best examples. Many sedimentary rocks and nearly all igneous and metamorphic rocks are composed of multiple minerals containing these two elements.

For example, titanium (Ti) may occur as a minor component in biotite, amphibole, or other minerals. In many rocks, however, Ti is concentrated in Ti-rich minerals such as rutile (TiO2), titanite (CaTiSiO5), and ilmenite (FeTiO3).

Similarly, rocks containing significant amounts of phosphorous usually contain apatite, Ca5(PO4)3(OH,F,Cl), or monazite, (Ce,La,Th,Y)PO4. Here (Figure 2.3) we see the Periodic Table of the Elements.

The first element is hydrogen (atomic number 1) and the last is oganesson (atomic number 118). Oganesson is an obscure synthetic radioactive element that has only been produced a few times and in minute amounts.

The chart contains 118 elements in all, but only about 90 occur naturally. The other 20 or so (with atomic numbers 104-118 and 99-103) elements are artificial – they are synthesized in nuclear reactors and some are radioactive with very short lives.

Although the Periodic Table of the Elements has appeared in many forms, the basic relationships are the same today as they were when Dmitri Mendeleyev (1834-1907) devised the first version in 1870. Some elements were unknown and omitted from the original table during Mendeleyev’s time, but were later discovered and added.

Chemists classify elements into different types that have related properties (shown by different colors in the table above). Hydrogen is a special element, but the other elements in Group 1 are alkali metals.

Those elements in Groups 4 through 12 are transition metals. Group 17 elements are the halogens, and Group 18 elements are the noble gases.

The asterisks in the chart show where the lanthinides (also called rare earth elements) and actinides (both considered transition metals) were extracted from the main chart and put as separate rows at the bottom. If we did not extract them, the chart would be too wide to fit easily on this page.

Elements in the same groups (columns) have their outermost electrons in the same kind of orbitals. This is very significant because it means the elements have similar chemical properties, form similarly charged ions and, importantly, commonly substitute for each other in mineral crystals.

When it first crystallized from a magma, it was a homogeneous mixture of Na-feldspar and K-feldspar. Na and K, both alkali elements (Group 1), commonly substitute for each other in minerals.

The result is called exsolution (which means unmixing) and the specimen now contains thin pinkish veins of K-feldspar surrounded by Na-feldspar. We call feldspar that has exsolved like this, perthite.

Z is also equal to the number of electrons orbiting the nucleus in neutral (non-ionized) atoms, and is close to the number of electrons in most ions. A neutral iron atom (Fe0), for example, contains 26 protons (Z = 26) in its nucleus and 26 electrons in an electron cloud around the nucleus.

Atomic nuclei (except one isotope of hydrogen) contain neutrons in addition to protons, and the number of neutrons, designated by N, may vary. This leads to isotopes of different mass numbers.

Most chemical elements have several different naturally occurring isotopes. Some isotope varieties, however, are generally more common than others, and some only exist in minute amounts.

Oxygen may be 16O, 17O, or 18O, where the superscript number denotes A. Examination of the equation in the previous paragraph tells us that the three isotopes of oxygen must have 8, 9, and 10 neutrons, respectively, because all must have 8 protons if they are oxygen.

16O is 99.8% of all natural oxygen. A mole of an element (or of a compound) is defined as containing 6.022 × 1023 atoms (or molecules).

So, one mole of carbon is equivalent to 6.022 × 1023 carbon atoms. The scale used to measure atomic mass has changed slightly over time.

Consequently, all atomic masses are given in atomic mass units (amus), defined as one-twelfth the mass of 12C. Both protons and neutrons have equivalent mass, about one amu, and electrons have almost no mass (less than 1/1000th the mass of protons and neutrons).

But for several reasons, not worth going into here, the mass of an atom is generally close to, but not the same as, the mass number. Elements are different from atoms.

Atomic masses/weights of elements are molal quantities (and often given in units of grams/mole) but they are really dimensionless numbers because they are all calculated relative to the atomic mass/weight of a mole of carbon. Although isotope mass numbers are always integers, atomic weights of elements are not.

When elements combine to produce a compound, the atomic weight of the compound is the sum of the weights of the elements in the compound. FeO, for example, has atomic weight 71.846 (15.999 + 55.847) grams/mole.

Thus, most quartz (SiO2) contains about the same relative amounts of the three natural oxygen isotopes (mostly 16O) depicted in Figure 2.5 above. Furthermore, isotopic variations have extremely small effects on the properties of minerals.

Small isotopic variations, however, may be significant to a geochemist trying to determine the genesis of a particular mineral or rock. Quartz (SiO2) is one of the most common and well-known minerals.

How much quartz is that.

By the end of this section, you will be able to: In ordinary chemical reactions, the nucleus of each atom (and thus the identity of the element) remains unchanged.

The transfer and sharing of electrons among atoms govern the chemistry of the elements. During the formation of some compounds, atoms gain or lose electrons, and form electrically charged particles called ions (Figure 1).

(a) A sodium atom (Na) has equal numbers of protons and electrons (11) and is uncharged. (b) A sodium cation (Na+) has lost an electron, so it has one more proton (11) than electrons (10), giving it an overall positive charge, signified by a superscripted plus sign.

Atoms of many main-group metals lose enough electrons to leave them with the same number of electrons as an atom of the preceding noble gas. To illustrate, an atom of an alkali metal (group 1) loses one electron and forms a cation with a 1+ charge.

For example, a neutral calcium atom, with 20 protons and 20 electrons, readily loses two electrons. This results in a cation with 20 protons, 18 electrons, and a 2+ charge.

The name of a metal ion is the same as the name of the metal atom from which it forms, so Ca2+ is called a calcium ion. When atoms of nonmetal elements form ions, they generally gain enough electrons to give them the same number of electrons as an atom of the next noble gas in the periodic table.

atoms of group 16 gain two electrons and form ions with a 2− charge, and so on. For example, the neutral bromine atom, with 35 protons and 35 electrons, can gain one electron to provide it with 36 electrons.

It has the same number of electrons as atoms of the next noble gas, krypton, and is symbolized Br−. (A discussion of the theory supporting the favored status of noble gas electron numbers reflected in these predictive rules for ion formation is provided in a later chapter of this text.).

Moving from the far left to the right on the periodic table, main-group elements tend to form cations with a charge equal to the group number. That is, group 1 elements form 1+ ions.

Moving from the far right to the left on the periodic table, elements often form anions with a negative charge equal to the number of groups moved left from the noble gases. For example, group 17 elements (one group left of the noble gases) form 1− ions.

This trend can be used as a guide in many cases, but its predictive value decreases when moving toward the center of the periodic table. In fact, transition metals and some other metals often exhibit variable charges that are not predictable by their location in the table.

Figure 2. Some elements exhibit a regular pattern of ionic charge when they form ions.

What is its symbol.

Knowing this lets us use the periodic table to identify the element as Al (aluminum). The Al atom has lost three electrons and thus has three more positive charges (13) than it has electrons (10).

Give the symbol and name for the ion with 34 protons and 36 electrons. Magnesium and nitrogen react to form an ionic compound.

Write the symbol for each ion and name them. Magnesium’s position in the periodic table (group 2) tells us that it is a metal.

A magnesium atom must lose two electrons to have the same number electrons as an atom of the previous noble gas, neon. Thus, a magnesium atom will form a cation with two fewer electrons than protons and a charge of 2+.

Nitrogen’s position in the periodic table (group 15) reveals that it is a nonmetal. Nonmetals form negative ions (anions).

Thus, a nitrogen atom will form an anion with three more electrons than protons and a charge of 3−. The symbol for the ion is N3−, and it is called a nitride ion.

Predict which forms an anion, which forms a cation, and the charges of each ion. Write the symbol for each ion and name them.

When electrons are transferred and ions form, ionic bonds result. Ionic bonds are electrostatic forces of attraction, that is, the attractive forces experienced between objects of opposite electrical charge (in this case, cations and anions).

Covalent bonds are the attractive forces between the positively charged nuclei of the bonded atoms and one or more pairs of electrons that are located between the atoms. Compounds are classified as ionic or molecular (covalent) on the basis of the bonds present in them.

The compound formed by this transfer is stabilized by the electrostatic attractions (ionic bonds) between the ions of opposite charge present in the compound.

A compound that contains ions and is held together by ionic bonds is called an ionic compound. The periodic table can help us recognize many of the compounds that are ionic: When a metal is combined with one or more nonmetals, the compound is usually ionic.

However, it is not always true (for example, aluminum chloride, AlCl3, is not ionic). You can often recognize ionic compounds because of their properties.

For example, sodium chloride melts at 801 °C and boils at 1413 °C. (As a comparison, the molecular compound water melts at 0 °C and boils at 100 °C.) In solid form, an ionic compound is not electrically conductive because its ions are unable to flow (“electricity” is the flow of charged particles).

Figure 3. Sodium chloride melts at 801 °C and conducts electricity when molten.

Watch this video to see a mixture of salts melt and conduct electricity.

These molecular compounds (covalent compounds) result when atoms share, rather than transfer (gain or lose), electrons. Covalent bonding is an important and extensive concept in chemistry, and it will be treated in considerable detail in a later chapter of this text.

Under normal conditions, molecular compounds often exist as gases, low-boiling liquids, and low-melting solids, although many important exceptions exist. Whereas ionic compounds are usually formed when a metal and a nonmetal combine, covalent compounds are usually formed by a combination of nonmetals.

While we can use the positions of a compound’s elements in the periodic table to predict whether it is ionic or covalent at this point in our study of chemistry, you should be aware that this is a very simplistic approach that does not account for a number of interesting exceptions. Shades of gray exist between ionic and molecular compounds, and you’ll learn more about those later.

Using the periodic table, predict whether the following compounds are ionic or covalent: Metals (particularly those in groups 1 and 2) tend to lose the number of electrons that would leave them with the same number of electrons as in the preceding noble gas in the periodic table.

Before you start this video, it is helpful to know how to name an ionic compound based on its chemical formula. In this lesson, you will learn how to start with the name of the ionic compound and turn it in to a chemical formula.

Binary ionic compounds are fairly simple. An ionic compound is a neutrally charged compound that is made up of bonded ions, a cation and an anion.

When the two combine into a compound, that compound does not have an overall charge. NaCl is an example.

Before you learn the steps for writing a chemical formula, I’d like to remind you how to determine the charge on an ion. For the representative elements, the charge of the ion is related to the column or group that the element is in.

Transition metals are elements are capable of losing different numbers of electrons and can take multiple ionic forms. The names of transition metal ions contain Roman numerals to indicate the ions’ charge.

For example, lead (II) nitrate contains a +2 lead ion: Pb2+. Vanadium (IV) oxide contains a +4 vanadium ion: V4+.

How big is an atom. A simple question maybe, but the answer is not at all straighforward.

However, even for atoms of the same type, atomic radii can differ, depending on the oxidation state, the type of bonding and – especially important in crystals – the local coordination environment. Take the humble carbon atom as an example: in most organic molecules a covalently-bonded carbon atom is around 1.5 Ångstroms in diameter (1 Ångstrom unit = 0.1 nanometres = 10-10 metres).

In the following article we’ll explore a number of different sets of distinct atomic radius sizes, and later we’ll see how you can make use of these “preset” values with CrystalMaker. Atomic radii represent the sizes of isolated, electrically-neutral atoms, unaffected by bonding topologies.

Atomic radii decrease, however, as one moves from left to right, across the Periodic Table. Although more electrons are being added to atoms, they are at similar distances to the nucleus.

Atomic radii are generally calculated, using self-consistent field functions. CrystalMaker uses Atomic radii data from two sources:

Springer Verlag, Berlin. CPK Atomic Radii: Clementi E, Raimondi DL, Reinhardt WP (1963).

The covalent radius of an atom can be determined by measuring bond lengths between pairs of covalently-bonded atoms: if the two atoms are of the same kind, then the covalent radius is simply one half of the bond length.

Van-der-Waals radii are determined from the contact distances between unbonded atoms in touching molecules or atoms. CrystalMaker uses Van-der-Waals Radii data from:

These are the “realistic” radii of atoms, measured from bond lengths in real crystals and molecules, and taking into account the fact that some atoms will be electrically charged. For example, the atomic-ionic radius of chlorine (Cl-) is larger than its atomic radius.

dAB = rA + rB. CrystalMaker uses Atomic-Ionic radii data from:

Perhaps the most authoritative and highly-respected set of atomic radii are the “Crystal” Radii published by Shannon and Prewitt (1969) – one of the most cited papers in all crystallography – with values later revised by Shannon (1976). These data, originally derived from studies of alkali halides, are appropriate for most inorganic structures, and provide the basis for CrystalMaker’s default Element Table.

Shannon RD Prewitt CT (1969) Acta Crystallographica B25:925-946. Shannon RD (1976) Acta Crystallographica A23:751-761.

Of course, atoms don’t have “colour” in the conventional sense, but various conventions have been established in different disciplines. Many organic chemists use the so-called CPK colour scheme These colours are derived from those of plastic spacefilling models developed by Corey, Pauling and (later improved on by) Kultun (“CPK”).

Organic Structures Alert. CrystalMaker’s default Element Table is the Shannon & Prewitt “Crystal” radii, which is appropriate for most inorganic structures.