Bonding and Chemical Formulas

Unit IVA - .2 Chemistry

October 2006- Revised Nov. 2007

Bonding

· There are two major types of bonding:· Ionic and covalent:· Covalent bonds occur when two atoms SHARE

valence electrons · Ionic bonds are due to a very strong

electrostatic attraction between two ions (atoms who just “EXCHANGED” electrons)

· In other words, one atom stole electrons from the other

Ions

· Let’s start by deciding what an ion is· An ion is an atom which has either

gained or lost an electron· When an atom gains an electron, the ion

is negatively charged· Why?· When an atom loses an electron, the ion

is positively charged· Why?

Ions

· A positively charged ion is referred to as a cation

· A negatively charged ion is called an anion

· Metals tend to be cations because they will frequently lose electrons

· Nonmetals tend to be anions because they will take on more electrons

Electronegativity

· The difference in electronegativities between two atoms, determines whether an ionic or covalent bond will form.

· So, what is electronegativity?

Electronegativity is defined as :

an atom’s pull on its electrons

As a rule of thumb, electronegativity increases as you go and to the

Electronegativity· Look at your table of electronegativities· Metals tend to have very LOW electronegativity

values and NonMetals tend to have quite HIGH values

· The difference between any two atoms will decide whether an ionic or a covalent bond will form

· If the diff is >2, then an ionic bond will form· If the diff is <1.69, then a covalent bond will form· (between 1.7-2, if M-NM then ionic and if 2 NM

then covalent)

Ionic vs Covalent

· What type of bond will form between each of the following?H and Cl?

· C and O?

· Na and Br?

· Be and F?

· What trend do you see?

· Typically nonmetals form ionic bonds with metals and they form covalent bonds with other nonmetals

Valence Electrons

· As you recall, the electrons in the outermost shell or energy level are called valence electrons

· These valence electrons are typically the only ones involved with bonding

· Let’s use a technique called electron dot structures to represent the valence electrons of an atom (also called Lewis Dot structures)

Valence Electrons· Each dot represents an electron· First, determine how many valence electrons an

atom has· Where will you find this information?· Then write the chemical symbol for that atom· Place the correct number of dots around each of

the 4 sides of the atomic symbol· Be sure to only put one dot per side until all

sides are “1/2 filled”, then you can start doubling up

· Why?

Electron Dot Structures

· Na

· C

· O

· Cl

· Kr

· Mg

· Li

· B

· P

· S

· Ne

· Ca

•Notice the patterns that will occur within families

•Write the electron dot structure for each of the following:

Electron Configurations for Ions

· We already know how to write the electron configuration for an atom, let’s apply it to ions

· We do it the exact same way except we take into account the electrons that have been added or the missing electrons

· The most stable form of the ion will be that which shares the electron configuration with a noble gas

Electron Configurations for Ions

· Noble gases all have the same ending to their electron configuration, s2 p6, giving them 8 electrons in their outermost energy level

· This is where the octet rule comes from· We know that noble gases are the most stable

elements, scientists gathered that the reason for this is because they have 8 electrons in their valence shell

Electron Configurations for Ions

· What is the noble gas configuration for the following ions?

· Ca+2

· O 2-

· Li+1

· Al+3

· H-1

Electron Configurations for Ions

· What is the noble gas configuration for the following ions?

· Ca+2 [Ar]· O 2- [Ne]· Li+1 [He]· Al+3 [Ne]· H-1 [He]

Ionic Bonding

· Ionic bonding involves the TRANSFER of electrons or stealing of electrons as we mentioned before

· Do you recall how we determine which element is going to become a cation and which will become an anion?

Ionic Bonding

· Each of you will be given a card with an elemental symbol on it

· For that symbol, determine the number of valence electrons (this is how many dots you would put around the symbol in your Lewis dot structure)

· Then, come up and grab $1 bill for each valence electron you have

Ionic Bonding

· Using the list of compounds, find your partner and EXCHANGE electrons accordingly

· Once you have exchanged electrons, you are transformed into your respective ions (cation/anion)

· You are electrostatically attracted to each other so… BOND (stand tightly shoulder to shoulder until I come around and check your compound)

Properties of Ionic Compounds

· A solid ionic compound is called a salt

· All salts share 5 characteristics:

· 1. Made of crystals

· 2. Conduct electricity

· 3. Have high melting and boiling points

· 4. Are hard

· 5. Are brittle

1. Made of Crystals

· Attraction between opposite charged ions is SO great that there is more than one bond

· A tightly packed cluster of repeating units forms the crystal structure

· Ex. NaCl is the formula· unit for table salt · crystals

2. Conducts Electricity

· Electricity needs charged particles that are free to move (in solution)

· In salt water, the particles spread out and can carry electricity from ion to ion throughout the solution

· Electrolytes - ions in solution which carry an electric current

3. High BP and MP

· Because of the strong attractions between the oppositely charged ions, it takes a lot of energy to “break up” the particles

· BP - boiling point is the temperature at which you have a phase change from a liquid to a gas

· MP - melting point is the temperature at which you have a phase change from a solid to a liquid

4. Hard and 5. Brittle

· Salts are hard due to the strong attraction of opposite charges and the layering of crystals

· Salts are brittle, or break up to make a powder

· Layers usually line up so that - and + alternate, but added energy or pressure can cause + to be next to + and - to be next to -.

· Does it like it…NO so it breaks apart into powder

Hydrates

· Some salts can hold water molecules between their bonds

· These are called Hydrated Salts· Possible Uses: drying agents or moisture

indicators· Ex. CuSO4 . 5 H2O· In this hydrate, for every salt unit of Copper

sulfate, 5 molecules of water are trapped

Percent Composition of Hydrates

· What % of CuSO4 . 5 H2O is water?

· First let’s find the formula mass of copper sulfate by itself

· 63.5 + 32.1 + 64.0 = 159.6 g/mol

· What about the water that is trapped?

· 1.0 (10) + 16.0 (5) = 90.0 g/mol

· Total molar mass: 249.6 g/mol

· % mass of water = 90.0 / 249.6 x 100 = 36%

Try another hydrate problem:

· Calculate the percent water in NiCl2. 6H2O ?

· NiCl2 = 129.7 g/mol

· 6 H2O = 108.0 g/mol

· Total mass: 237.7 g/mol

· 108.0 / 237.7 x 100 = 45%

Percent Composition

· The law of definite proportions refers to the chemical make-up of ONE compound

· Within that compound, the proportion or ratio of one element to another will remain the same no matter how much of the compound is present

· The law of multiple proportions compares the compositions of two different compounds which contain the same elements

Percent Composition

· The Law of definite proportions explains why we can have a “formula unit” for a compound

· This is the simplified version of the elemental ratios within the compound

· For example, when joining Mg and Cl what is the smallest whole number ratio that can be used to join these two together

· Make sure that the charges balance out

Percent Composition - (refresher)

· If I have an ionic compound of MgCl2, what percentage of the whole mass does Mg make up?

· First, find the atomic mass of Mg· Then, find the atomic mass of Cl and

double it because there are two atoms of Cl· What is the total mass?· Divide the mass of the Mg by the total mass

and multiply by 100. This is the percent of the whole that Mg makes up.

Percent Composition

· Calculate the percent composition of water in the following hydrate:

· (NH4)2SO4 . 5 H2O

Mole Ratios and Hydrate Predictions

· For our hydrated salt lab, you took the mass of the hydrated salt before heating

· You then heated it up for 10-15 minutes as instructed and took the mass of the anhydrous salt

· Anhydrous salt – salt without the water· To determine the amount of water that was

in your original hydrated salt sample, subtract the mass of anhydrous salt from the mass of hydrated salt

· Now you can do two things:

· 1. determine the percent of the salt sample that was water

· Do this by taking the mass of water and dividing by the total mass of the hydrated salt (before heating)

· Multiply by 100 and this is your experimental % water.

· 2. The second thing you can calculate is the # of waters trapped in the salt per formula mass unit.

· Do this by first converting your number of grams of anhydrous salt to moles (use the molar mass of the salt)

· Then convert your mass of water to moles (using the molar mass of water)

· Now divide both numbers (of moles) by whichever is smaller to get a ratio.

· Clue: the ratio will by 1 to ___ (rounded to the nearest whole number).

· If the ratio is very close to a HALF number, then double both. Ex. 1 : 2.5 gets doubled to 2:5 ratio

Hydrate Prediction Example

· In the lab, you measure out 5.25g of barium chloride (BaCl2 )

· After heating, the mass is 4.50g.

· Calculate the amound of water that was in the hydrated salt.

· 5.25g – 4.50g =

· 0.75 g

· Calculate the % of water in the sample:

· 0.75g / 5.25 g = 14.3% water

· Now, let’s determine the # of water molecules trapped in the hydrated salt per formula unit.

· Convert the grams of water to moles.

· 0.75g / 18.01 g/mol = 0.0416 moles H2O

· Convert the grams of anhydrous salt to moles.

· 4.50g / 208.24g/mol = 0.0216 moles BaCl2

· Divide each by the lesser of the two

· 0.0416/0.0216 = 1.93 (round to 2)

· 0.0216/0.0216 = 1

· Write your whole number ratio.

· 1 BaCl2 : 2 H2O

· Finally, write your formula for the hydrated salt

· BaCl2 . 2 H2O

· To determine how close you were (or to calculate % error): calculate your theoretical percent water in the above formula as usual

· 2(18.01) / ( 208.24 + 2(18.01) )· Mass of water / total mass =

· 14.75%· Compare with your original 14.30% experimental

percent water: 14.75 – 14.30 = 0.45· % error is this difference divided by the theoretical:· 0.45 / 14.75 =6.6% error (not too bad for a first

year chemistry lab student)

Polyatomic Ions

· As you saw when we discussed and calculated oxidation numbers, you sometimes will see a “cluster” of atoms that have a combined overall charge

· These are called poly (many) atomic (atoms) ions (with a charge)

· Ex. MnO4 -1 the permanganate ion

· I have provided you with a list of polyatomic ions that you’re responsible for memorizing

Naming Ionic Compounds

1. When naming ionic compounds, begin with the cation.

This can either be a metal from the periodic table OR it could be a polyatomic cation

· What are some cations?

· Family I, II, or III, transition metals, or ammonium

Ionic Compounds· 1.When writing the name, start with the cation

· 2.Follow it by the anion (for now we’ll start with polyatomic ions that have a known charge)

· Choose a polyatomic anion that would balance with Cu+2

· Let’s use sulfate for this example

· Write the formula for this compound

· Now write the name of it

· CuSO4

Using Roman Numerals

· When the cation can vary in its oxidation number/ charge, we must use a Roman numeral to indicate its charge

· Ex. Iron can vary in its ox #

· Write the formula for iron bonded to sulfate IF the iron has a +3 charge

· Fe2(SO4)3

Using Roman Numerals

Fe2(SO4)3

When we name this compound, we must use a Roman numeral to indicate iron’s charge so the name of this compound is

Iron (III) sulfate because the iron has a +3 charge

Do NOT confuse this with the subscript, which indicates only how many atoms of iron we have

Ionic Compounds

· Do we have to use the Roman numeral for family I and II elements?

· No, they have a set charge that is understood· Practice writing the names of the following:

· ZnSO4 CaCO3

Fe2(SO4)3

· Zinc sulfate Calcium Carbonate Iron (III) Sulfate

Other Anions

· Polyatomic ions do not all end in ide

· but other ions do

· When naming, you still place the name of the cation first

· Followed by the anion (ending in –ide)

· What would you call MgO ?

· Magnesium oxide

Ionic Compounds (cont.)

· Name each of the following:· NaCl KI FrBr· Sodium chloride Potassium iodide Francium bromide

· What happens when you have multiple anions with a cation? For example: MgCl2 ?

· This is still called Magnesium chloride because in order to form this ionic compound, there HAS to be 2 chlorines