Chemistry 1, Chapter 7

Ions and Ionic Compounds

Section 1, Simple Ions

1. Chemical Reactivity

• Atoms react to achieve a stable electron configurations

– Remember, the stable arrangement is a full outer level of electrons, usually 8

• How reactive an element is depends on its outer electron configuration

– noble gases are non-reactive because they already have a full outer level

• The idea that 8 outer electrons (electrons in the s and p sublevels) is a full level and they make an atom stable or unreactive is called the octet rule

• There are threes ways that atoms can get a full outer level

1. they can gain electrons

2. they can lose electrons and expose a full level below

3. they can share electrons

• in this chapter we are going to learn about gaining and losing electrons

• From these three ways to get a full outer level we can form three types of bonds

1. ionic – which involves the gain and loss of electrons (between metal and nonmetal – chapter 7)

2. covalent – which involves the sharing of electrons (between nonmetals – chapter 8 )

3. metallic bonds- which involves sharing of electrons (between metals - chapter 7)

• Metals tend to lose electrons in reactions

– Remember, metals are usually found in groups 1-12, some in 13-16, and the lanthinides and actinides

– metals usually have 3 or fewer outer electrons and it is therefore easier for them to lose electrons than to gain them

– alkali metals are the most reactive metals

• Nonmetals tend to gain electrons in reactions

– remember, nonmetals are usually in groups 17 and 18, with some in groups 13-16

– nonmetals tend to have 5 or more outer electrons and it is therefore easier for them to gain electrons than to lose them

– halogens are the most reactive nonmetals

• Valence Electrons

– atoms of many elements become stable by achieving the outer electron arrangements of a noble gas

– remember, we call these outer electrons valence electrons

– we can determine an atoms valence electrons from the pattern on the periodic table, and from that we can predict whether the atom will lose or gain electrons

• group 1 – 1 outer electron (lose 1)• group 2 - 2 outer electrons (lose 2)• group 13 – 3 outer electrons (lose 3)• group 14 – 4 outer electrons (lose or gain 4)• group 15 – 5 outer electrons (gain 3)• group 16 – 6 outer electrons (gain 2)• group 17 – 7 outer electrons (gain 1)• group 18 – 8 outer electrons (already full, so they are non-

reactive)• group 3-12, transition metals can have different

configurations but we predict 2 outer electrons for them (lose 2) except for silver and gold for which we predict 1 (lose 1) unless we are told otherwise

– When atoms gain or lose electrons they are making their number of protons and electrons unequal

• this means the atoms will no longer remain neutral

• when atoms are not neutral they must have a charge and we call charged atoms ions

• if an atom loses electrons, it will have more protons (+) than electrons (-) and its charge will be positive

• we call positive ions “cations”

• their positive charge is equal to the number of electrons they lose

• if an atom loses 1 electron it will have a +1 charge, if it loses 2 electrons it will have a +2 charge and so on

• if an atom gains electrons, it will have more electrons (-) than protons (+) and its charge will be negative

• we call negative ions “anions”

• their negative charge will be equal to the number of electrons they gain

• if an atom gains 1 electron its charge will be –1, if it gains 2 electrons its charge will be –2, and so on

• so we can use the number of valence electrons to predict how many electrons an atom will lose or gain and therefore the charge it will take on- group 1 – 1 outer electron (lose 1) = +1– group 2 - 2 outer electrons (lose 2) = +2– group 13 – 3 outer electrons (lose 3) = +3– group 14 – 4 outer electrons (lose or gain 4) = + or - 4– group 15 – 5 outer electrons (gain 3) = -3– group 16 – 6 outer electrons (gain 2) = -2– group 17 – 7 outer electrons (gain 1) = -1– group 18 – 8 outer electrons (already full, so they are

non-reactive)– group 3-12, transition metals can have different

configurations but we predict 2 outer electrons for them (lose 2 = +2) except for silver and gold for which we predict 1 (lose 1 = +1) unless we are told otherwise» transition metals can often form more than one ion, we predict the

ions as mentioned above unless we are told otherwise

Predict +2, except Au & Ag = +1

↓ ↓

• therefore, we can see that metals form cations (+) by losing electrons and nonmetals form anions (-) by gaining electrons

• notice the nucleus is never changed in this process

• label and learn the pattern for prediction ionic charge on the periodic table and your flashcards

• Atoms and Ions– When atoms lose and gain electrons to form ions,

they are still the same elements• remember, the number of protons is what determines

the identity of an atom• just because an atom is charged, it still has the same

number of protons and therefore the same identity

– Ions have different properties than atoms that made them

– Ions are different sized than atoms that made them

• atoms that lose electrons form ions that are smaller than their parents

• atoms that gain electrons form ions that are larger than their parents

• Because opposites attract, cations and anions attract one another

– This is what happens when ionic bonds form to make ionic bonds

• the word salt actually means an ionic compound that forms when a metal atom or a positive radical (group of two or more atoms acting as a single atom) replaces the hydrogen of an acid

• examples of salts: sodium chloride, potassium chloride, magnesium oxide etc.

• all of these salts are ionic compounds that are electrically nuetral

• anions actually attract several cations and vice-versa, in this way many ions are pulled together into a tightly packed structure

•this structure gives any salt its crystalline structure

•the smallest crystal of table salt that you can see still has more than a billion billion sodium and chloride ions

• remember ionization energy is the energy required to remove an electron and electron affinity is the energy required to add an electron onto a neutral atom

• this is only a part of the process involved in forming an ionic bond

• adding and removing electrons requires energy

• more energy is released when ionic bonds are formed than is required to make the ions

• since energy is released the overall process is exothermic and spontaneous, even though some parts require energy to occur

• also, since energy is released when these bonds are formed, then energy is required to break these bonds

Ionic Compounds

• remember that cations are positive and anions are negative, but when they come together to form an ionic bond they have no overall charge– the ratio of the charges cancel each other out

• ionic compounds do not consist of molecules – remember that ionic compounds are formed

when anions and cations attract each other to form a crystal

• elements in group 1 and 2 reacting with groups 16 and 17 will almost always form ionic compounds and not molecular (covalent) compounds

• ionic bonds are very strong which gives ionic compounds certain properties– most ionic compounds have high melting and boiling points,

and they are usually solid at room temperature– liquid and dissolved salts conduct electrical currents

• to conduct an electric current, a substance must satisfy 2 conditions– they must contain charged particles

– those particles must be free to move

• ionic solids, such as salts, do not usually conduct electrical current because the ions are not free to move

• however, when they melt or dissolve their ions are free to move and are therefore excellent conductors

• there is a small class of ionic compounds that can allow charges to move through their crystals

– ex. Zirconium oxide

– Salts are hard and brittle

– how to identify ionic compounds (characteristics)• solid at room temperature• hard and brittle• high melting and boiling points• good conductors of electrical current in the liquid

state (melted) or dissolved in water

• Salt Crystals– the ions in salts form repeating patterns

• not all salts have the same crystal structure

• the crystals of all salts are made of simple repeating units, called a crystal lattice

– crystal structures depend on the sizes and ratios of ions

– Salts have ordered packing arrangements

– the smallest repeating unit in a crystal lattice pattern is called a unit cell

Names and Formulas of Ionic Compounds

• Naming Ionic Compounds

– naming salts (ionic compounds) is very easy

• simple ionic compounds made of just two elements are called binary compounds, we will learn about these first

• rules for naming simple ions– cations use the same name as their parent atom

• ex.: K+ is the potassium ion, Zn+2 is the zinc ion

– when elements, such as transition metals, form more than one type of ion, the ion name includes a roman numeral to indicate its charge

• ex.: Copper has 2 common ions, one with a +1 cahrge and one with a +2 charge; so, we name them copper (I) and copper (II) with symbols Cu(I) and Cu (II)

– anions also get names from their elements, but we change their endings to “ide”

• examples: Chlorine forms an anion with a -1 charge, its symbol is Cl – and we call it a chloride ion. Oxygen forms an ion with a -2 charge, its symbol is O-2 and we call it an oxide ion.

• to name ionic compounds we use the names of the ions present in the compound

– we write the name of the cation first followed by the anion

– example: NaCl is sodium chloride. ZnS is zinc sulfide. K2O is potassium oxide. CuCl2 is copper (II) chloride. Al2S3 is aluminum sulfide.

• Writing Ionic Formulas

– remember, compounds have no overall charge

• this means that the charges from the cations and anions must cancel each other out mathematically

• steps for writing formulas of binary compounds– write the symbol and charges for the cation and anion

– find the lowest common multiple of the charges to see how many of each ion you need to cancel the charges

– example: magnesium nitride is made of Mg+2 and N-3 , we need to figure out how many of each ion we need to get the charges to cancel. Find the least common multiple of 3 and 2 which is six, so we need to have 6 positive charges and 6 negative charges to balance out. How can we get 6 positives; each Mg has 2 positives, so we need 3 of them. How can we get 6 negatives; each N has 3 negatives, so we need 2 of them. We then write the formula to reflect how many of each ion we have. Mg3N2 . Notice that we no longer write the charges, and the numbers of each ion are written as subscripts.

• remember, ionic compounds are made of crystal lattices (not molecules), therefore, their formulas show us the smallest ratio of charges needed to be neutral

• there is a shortcut to writing ionic formulas; the cross-over method – we take the charge from the cation and write it as the

subscript of the anion and vice-versa– If the subscripts can be reduced mathematically, reduce

them

– Ex.1 Mg+2 and Cl-1 , MgCl2

– Ex. 2 Mg+2 and O-2 , Mg2O2, MgO

Polyatomic Ions • Up till now, we have been talking about simple

ions. This means ions made from one charged atom. – we also call these monatomic ions

• But there are ions that are made of more than one atom. These are called polyatomic ions.– polyatomic ions are a charged group of bonded atoms

that act together as one atom – they usually act the same way simple ions do– they can either be cations or anions– common polyatomic ions, you will have to memorize

these

We call the charges on the ions oxidation numbers.

• the names of polyatomic ions can be complicated– many polyatomic ions contain oxygen; endings such as “ite”

and “ate” indicate the presence of oxygen

– many polyatomic ions differ only by the number of oxygen atoms present

• if there are only 2 forms; the form with less oxygen ends in “ite”, the form with more oxygen ends in “ate”

• if there are more than 2 forms; we can use the prefix “hypo” for the least number of oxygen and “per” for the most

• examples: nitrate (NO3-) and nitrite (NO2

-) ,or

• hypochlorite (ClO-) and chlorite (ClO2-) and chlorate (ClO3

-) and perchlorate (ClO4

-)

– the prefix “thio” means replaced an oxygen with a sulfur• example: sulfate (SO4

-) and thiosulfate (S203-)

• Naming compounds with polyatomic ions in them

– we name them the same as binary compounds except if the compound ends in the polyatomic ion (anion) we do not use and “ide” ending, but rather keep the ending of the polyatomic ion

– example: NH4Cl is ammonium chloride and CaCO3 is calcium carbonate

• Writing a formula for a compound containing a polyatomic ion

– we also write formulas for compounds with polyatomic ions the same as we do for binary compounds

– the only thing we need to remember is that the polyatomic ion acts as one atom, so if we need to add a subscript to a polyatomic ion we must put parentheses around it to show that the subscript goes with all of its atoms

– If you don’t use parentheses, you will be changing the formula and therefore the compound

• example: to write the formula for magnesium hydroxide (Mg+2 and OH-), we can see that we need 2 OH’s to cancel the magnesium. When we write the formula, we put parentheses around the hydroxide and then add the subscript so we keep the identities of the ions the same

– Mg(OH)2 is correct, if we do not include the parentheses we get MgOH2 which is incorrect