Unit 4

Formulas and Equations

Textbook Chapter 2, 6, & 8

Review Book Topic 2

Chemical Symbols

• Each element has been assigned a one-, two- or three- letter symbol for its identification

• First letter is ALWAYS capitalized, additional letters are lowercase

• Only recently discovered, unnamed elements are given three- letter symbols

• Some symbols show a relationship

– Ex. Carbon ~ C

Sodium ~ Na (Latin – natrium)

• Symbols are assigned by IUPAC

– International Union of Pure and Applied Chemists

• Roots used for naming elements:

0 : nil 1 : un 2 : bi 3: tri 4 : quad

5 : pent 6 : hex 7 : sept 8 : oct 9 : enn

• Ex. Element #109

– Un-nil-enn-ium

(1)-(0)-(9)

• Ex. Element #114

Chemical Molecules

• Monatomic molecules – uncombined elements, written without a subscript

– Ex. Neon gas – Ne

Argon gas – Ar

• Diatomic molecules – elements can exist in nature as two identical atoms bonded together

– Ex. Hydrogen – H2

(F, O, N, Cl, Br, I)

Chemical Formulas

• Chemists have identified over 10 million compounds

• Compound – two or more elements that are chemically combined (bonded together) in definite proportions by mass

– Ex. H2O, C6H12O6, H2O2

• No two compounds have identical properties

• Formulas use chemical symbols and numbers to show what elements and how many atoms of each are involved in each compound

• Chemical formula – shows the kinds and numbers of atoms in the smallest representative unit of the substance

– If monatomic: use chemical symbol (ex. Kr)

– If diatomic or a compound: use chemical symbols of elements involved, and subscripts to represent # of atoms present (ex. F2 or O3 or NaCl)

– Types of formulas: molecular, empirical, structural

• Subscript – smaller number after an element symbol that indicates how many atoms of that element are in the molecule

– Ex. H2O means there are 2 H and 1 O atom

• Coefficient – number in front of a molecule’s formula indicating how many molecules are present

– Ex. 2H2O means there are 2 water molecules

• Molecular formulas – shows the kinds and numbers of atoms present in a molecule of a compound

– Subscript written after the symbol indicates the # of atoms of each element

• If only 1 atom, subscript of 1 is omitted

– Show composition but NOT molecular structure

• Empirical formula (“formula unit”) – shows the lowest whole number ratio of ions in a compound

– Ex. MgCl2

• For every 1 Mg+, there are 2 Cl-

– Ex. H2O and H4O2

• Both have a ratio of 2 H : 1 O

• Molecular formulas can be seen as a multiple of an empirical formula

– Ex. Glucose: C6H12O6 (molecular)

CH2O (empirical)

6(CH20) = C6H12O6

• Structural formula – shows the physical organization of the atoms in a molecule

• Law of definite proportions – in any compound, the masses of the elements involved are always in the same proportions

– Ex. NaCl always has 1 Na (23 amu) and 1 Cl (35 amu) = 58 amu total for one NaCl

– Ex. H2O always has 2 H (total 2 amu) and 1 O (16 amu) = 18 amu total for one H2O

– Proportions of mass equals the ratio proportions of the number of atoms of each element in the molecule

• Law of multiple proportions – whenever two elements form more than one compound (ex. H2O and H2O2), the different masses of one element (ex. O versus O2) that combine with the same mass of the other element (H2) are in the ratio of small whole numbers

– Ex. We have two compounds, each with 2 g of element B. Compound 1 has 5g element A, compound 2 has 10 g element A

Review Questions

1. Determine the empirical formula for the compounds below:

a. C6H8O6 (vitamin c)

b. Hg2Br2

c. K2CrO4

d. N2H4

e. C6H6

f. C2H4O2

Review Questions

1. Determine the empirical formula for the compounds below:

a. C6H8O6 (vitamin c) C3H4O3

b. Hg2Br2 HgBr

c. K2CrO4 K2CrO4

d. N2H4 NH2

e. C6H6 CH

f. C2H4O2 CH2O

Atoms, Compounds and Ions

• Atoms and compounds are electrically neutral

– (# p+ = # e-)

• Ions have a net charge, either (+) or (-)

– (# p+ ≠ # e-)

– (+) ions attract (-) ions in a ratio that produces a neutral compound

Monatomic Ions

• Ions consisting of only one atom

• Ionic charges are found using the periodic table (look at group #s)

• Metals have a (+) ionic charge, nonmetals have a (-) ionic charge

• Metallic elements tend to lose electrons (forming cations)

– Group 1: 1+ charge

– Group 2: 2+ charge

– Aluminum: 3+ charge

• Nonmetallic elements tend to gain electrons (forming anions)

– N, P, and As: 3- charge

– O, S and Se: 2- charge

– F, Cl, Br, I (group 17): 1- charge

– Nonmetal ionic charge is found by subtracting the group number (in the form of 5A, 6A, 7A, etc.) from 8

• Group 0 usually does not form ions (noble gases)

• Transition metals tend to have more than one ionic charge (represented by the oxidation numbers on the periodic table)

• Determining the ionic charge of transition metals:

– Roman numerals are used in parentheses to indicate the numerical charge

– Form: Element name(roman numeral) ion

• No spaces are used between the element name and the first parentheses

• Ex. Cu 2+: Copper(II) ion

Sn 4+: Tin(IV) ion

• Roman Numerals:

Polyatomic Ions

• Tightly bound group of atoms that behave as a unit

• Carry an overall charge (+ or -)

• Reference Table E

• Names usually end in –ite or –ate

• Three exceptions:

– Ammonium cation (NH4+)

– Two polyatomic anions ending in –ide• Cyanide (CN-)• Hydroxide ion (OH-)

• -ite/-ate pairs of polyatomic ions:

-ite -ate

SO32-, sulfite SO4

2-, sulfate

NO2-, nitrite NO3

-, nitrate

ClO2-, chlorite ClO3

-, chlorate

• The –ite ending indicates one less oxygen atom than the –ate ending

• When the formula for a polyatomic ion begins with a hydrogen ion (H+):

– The charge on the new ion is the sum of the ionic charges:

• H+ + CO32- HCO3

- (hydrogen carbonate)

• H+ + PO43- HPO4

2- (hydrogen phosphate)

• H+ + HPO42- H2PO4

- (dihydrogen phosphate)

Review Questions

1. How can the periodic table be used to determine the charge of an ion?

2. Explain what is meant by a polyatomic ion

Review Questions

1. How can the periodic table be used to determine the charge of an ion?

Look up the oxidation #s to see the different ion possibilities for each element or look at the group number

2. Explain what is meant by a polyatomic ion

Review Questions

1. How can the periodic table be used to determine the charge of an ion?

Look up the oxidation #s to see the different ion possibilities for each element or look at the group number

2. Explain what is meant by a polyatomic ionContains more than one ion but acts as a single “unit” by carrying an overall charge

3. Using only your periodic table, write the formula for the typical ion of each element and identify it as an anion or cation:

a. Potassium

b. Sulfur

c. Argon

d. Bromine

e. Beryllium

f. Sodium

3. Using only your periodic table, write the formula for the typical ion of each element and identify it as an anion or cation:

a. Potassium K+ cation

b. Sulfur S2- anion

c. Argon NONE

d. Bromine Br1- anion

e. Beryllium Be2+ cation

f. Sodium Na+ cation

• Write the formula, including the charge, for each ion:

a. Ammonium ion

b. Tin(II) ion

c. Chromate

d. Nitrate ion

e. Cyanide ion

f. Iron(III) ion

g. Permanganate ion

h. Manganese(II) ion

• Write the formula, including the charge, for each ion:

a. Ammonium ion NH4+

b. Tin(II) ion Sn2+

c. Chromate CrO42-

d. Nitrate ion NO3-

e. Cyanide ion CN-

f. Iron(III) ion Fe3+

g. Permanganate ion MnO4-

h. Manganese(II) ion Mn2+

Ionic Compounds

• Monatomic or polyatomic ions attract each other in a ratio that produces a neutral compound

– Opposite charges attract !

– Ex. HCl (H+ + Cl-)

– Ex. H2SO4 (H22+ + SO4

2-)

– Ex. AgNO3 (Ag+ + NO3-)

• A compound’s name should indicate its composition, behavior, and how it is related to other compounds….common names do not tell us anything about chemical composition!

– Ex. Sodium chloride versus salt

– Ex. Dihydrogen oxide versus water

• Binary compounds – composed of two elements

– (+) charge of the cation must balance with the (-) charge of the anion (equal but opposite charges)

– Net ionic charge = 0

• Formulas for ionic compounds are usually written with the cation first, followed by the anion and ALWAYS show the lowest whole-number ration of ions in the compound

• Ternary ionic compounds – composed of atoms of three different elements

– Usually contains a polyatomic ion

– Parentheses can be used around the polyatomic ion to show if more than one are used in a reaction

Writing Ionic Compound Formulas

• Crisscross method:

– The charge of each ion is crossed over and used as a subscript for the other ion

• For many elements, the oxidation state is equal to the charge on the ion

– The signs (+ or -) are dropped when used as subscripts

Review Questions

• Create a molecular formula between the following ions:

1. Fe3+ O2-

2. Ca2+ S2-

3. Ba2+ S2-

4. Li+ O2-

5. Ca2+ N3-

6. Cu2+ I-

7. K+ N3-

Review Questions

• Create a molecular formula between the following ions:

1. Fe3+ O2- Fe2O3

2. Ca2+ S2- Ca2S2 = CaS

3. Ba2+ S2- Ba2S2 = BaS

4. Li+ O2- Li2O

5. Ca2+ N3- Ca3N2

6. Cu2+ I- CuI2

7. K+ N3- K3N

• Ex. Na+ Cl-

• Ex. Mg2+ Cl-

• Ex. Ca2+ NO3-

• Ex. Na+ NO3-

• Ex. Na+ Cl- NaCl

• Ex. Mg2+ Cl- MgCl2

• Ex. Ca2+ NO3- Ca(NO3)2

• Ex. Na+ NO3- NaNO3

Naming Ionic Compounds

• Compounds are named according to the types of elements that form them

• Ionic compounds are named in a different manner than covalent compounds

• Metal ion is usually named first, with the anion (monatomic or polyatomic) last

• If more than one ion can be formed for a metal ion, show the ionic charge using roman numerals in parentheses

• Binary Ionic Compounds:

– Name of the (+) ion is

named as is

• Ex. Sodium is sodium in NaCl

– Name of the (-) ion is slightly changed to end in –ide

• Ex. Chlorine becomes chloride in NaCl

Review Questions

• Name the following compounds:

1. KCl

2. MgS

3. AlN

4. MgCl2

Review Questions

• Name the following compounds:

1. KCl Potassium chloride

2. MgS Magnesium sulfide

3. AlN Aluminum nitride

4. MgCl2 Magnesium chloride

• Ternary Ionic Compounds:

– When (+) ion is a metal:

• Use unmodified metal name plus the name of the negative polyatomic ion

• Ex. KNO3 is potassium nitrate

– Most polyatomic ions are (-) charged

• Ammonium is an exception – NH4+

• Compounds containing this polyatomic ion are named differently:

– If (-) ion is a nonmetal, use the ending –ide

» Ex. NH4Cl is ammonium chloride

– If (-) ion is another polyatomic ion, each retain their names

» Ex. NH4NO3 is ammonium nitrate

Molecular Compounds

• Atoms of different elements combine to form compounds

• Molecular compounds are composed of molecules (smallest electrically neutral unit of a substance)

– Ex. H2O or CO2

• Characteristics:

– Low melting and boiling points

– Mostly liquids or gases at STP

– Composed of two or more non-metals

Characteristics of Molecular and Ionic Compounds

Characteristic Molecular Ionic

Representative Unit

Molecule Formula Unit (balance of oppositely charged

ions)

Type of Elements Nonmetals Metal combined with nonmetals

Physical State Solid, liquid, gas Solid

Melting Point Low (usually below 300⁰C)

High (usually above 300⁰C)

Naming Molecular Compounds

• Use a prefix system

– Tells how many atoms of each element are present in the compound/formula

– Prefixes represent the formula’s subscripts

• Names of all binary molecular compounds end in -ide

Prefixes Used in Naming Binary Molecular Compounds

Prefix Number

Mono- 1

Di- 2

Tri- 3

Tetra- 4

Penta- 5

Hexa- 6

Hepta- 7

Octa- 8

Nona- 9

Deca- 10

• Note:

– The vowel at the end of a prefix is dropped when the name of the element begins with a vowel

• Ex. Carbon monoxide not monooxide

– If there is a single atom of the first element name, the prefix mono- is dropped

Review Questions

• N2O

• PCl3

• SF6

• OF2

• Cl2O8

• SO3

Review Questions

• N2O Dinitrogen monoxide

• PCl3 Phosphorus trichloride

• SF6 Sulfur hexafluoride

• OF2 Oxygen difluoride

• Cl2O8 Dichlorine octoxide

• SO3 Sulfur trioxide

• Nitrogen trifluoride

• Disulfur dichloride

• Dinitrogen tetraoxide

• Tetraiodine nonoxide

• Phosphorus pentabromide

• Nitrogen trifluoride NF3

• Disulfur dichloride S2Cl2

• Dinitrogen tetraoxide N2O4

• Tetraiodine nonoxide I4O9

• Phosphorus pentabromide PBr5

Naming Acids

• Acids are compounds which produce hydrogen ions when dissolved in water

• Consider acids as combinations of anions connected to as many hydrogen ions (H+) as needed to make the molecule electrically neutral

• Memorize these:

– Hydrochloric acid HCl

– Sulfuric acid H2SO4

– Nitric acid HNO3

– Acetic acid HC2H3O2

– Phosphoric acid H3PO4

– Carbonic acid H2CO3

TABLE K IN REFERENCE TABLES

Naming Hydrates

• Hydrate – a compound with a specific amount of water bonded to it

– % of water = mass of water x 100%

mass of hydrate

• Anhydrous Salt – hydrates with water no longer attached

– Hydrate + Heat Anhydrate + Water

• How to name hydrates:

– Name the compound like normal

– Name the % of water by using prefixes

– Ex. CuSO4 * 5 H2O

– Ex. Na2CO3 * 10 H2O

• How to name hydrates:

– Name the compound like normal

– Name the % of water by using prefixes

– Ex. CuSO4 * 5 H2O

Copper (II) sulfate pentahydrate

– Ex. Na2CO3 * 10 H2O

Sodium carbonate decahydrate

Allotropes

• One of two or more different molecular forms on an element in the same physical state

– Ex. Oxygen (O2) and ozone (O3)

– Ex. Diamonds and graphite (both carbon)

• Allotropes have very different properties compared to each other even though the use the same element

Changing Properties

• Chemical property – ability of a substance to undergo a chemical reaction to form new substances

– Always results in a change in chemical composition

– Bonds are broken between atoms, atoms are rearranged, bonds are made to form new substances

– Irreversible

– Ex. Burning, rusting, decomposition, corrosion

• Physical property – quality or condition of a substance that can be observed or measured without changing the substances composition

– Includes phase changes or mixing of materials

– Reversible

– Ex. Cutting, grinding, bending

Types of Chemical Reactions

• Chemical reaction – one or more substances change into new substances

• Energy is absorbed or

given off

• Irreversible chemical

change

• Reactants Products

– Reactants: starting substances (left of arrow)

– Products: formed substances (right of arrow)

• Symbols are used to indicate the physical state of the substance in the equation

– s = solid

– l = liquid

– g = gas

– aq = aqueous (dissolved in water)

• Catalyst – a substance that speeds up the rate of a chemical reaction but is not used up in the reaction

– NOT a reactant OR a product

– Formula of the catalyst is written above the arrow if it is used in the reaction

Symbols Used in Chemical EquationsSymbol Explanation

+ Used to separate two reactants or two products

“Yields” - Used to separate reactants from products

↔ Used in place of in reversible reactions

(s) Designates a reactant or product in the sold state; placed after the formula

(l) Designates a reactant or product in the liquid state; placed after the formula

(g) Designates a reactant or product in the gaseous state; placed after the formula

(aq) Designates an aqueous solution; the substance is dissolved in water; placed after the formula

∆ Indicates that heat is supplied to the reactions

A formula written above or below the yield sign indicates its use as a catalyst

Types of Reactions

• Synthesis

– A + B AB

– 2 reactants 1 product

• Decomposition

– Single reactant is broken down into multiple products

– Reverse of synthesis

– AB A + B

• Single Replacement

– Reactants are 1 element + 1 compound

– A + BC AB + C

– Compounds are usually ionic

– Only occurs if…

• The more active (non)metal replaces the less active (non)metal

– Table J in reference tables

– Higher they are on the list, the more active they are

• Ex. Does the following reaction occur?

– 2 AgCl + Mg MgCl2 + 2 Ag

• Double Replacement

– AB + CD AD + BC

– Both reactants and products are compounds

– Double switch (must end with (+) and (–) ion in each compound)

– Occurs if…

• H2O is a product

• A gas is a product

• Precipitate is a product (insoluble) – Table F

–Table F is kinda weird, it is insoluble if it is in the INNER two columns

Review Questions

• CO2 + H2O H2CO3

• 3Ag + FeCl2 3AgCl + Fe

• 2Na + Cl2 2NaCl

Review Questions

• CO2 + H2O H2CO3

– Synthesis

• 3Ag + FeCl2 3AgCl + Fe– Single replacement– Will NOT occur – Ag is below Fe, cannot replace it

• 2Na + Cl2 2NaCl– Synthesis

• AgNO3 + NaCl AgCl + NaNO3

• 2H2O 2H2 + O2

• K2SO4 + CaCl2 CaSO4 + 2KCl

• AgNO3 + NaCl AgCl + NaNO3

– Double replacement– Will occur (a precipitate, AgCl, is formed)

• 2H2O 2H2 + O2

– Decomposition

• K2SO4 + CaCl2 CaSO4 + 2KCl– Double replacement– Will occur – calcium sulfate is a precipitate

Balancing Equations

• Law of Conservation of Mass – in any physical or chemical reaction, mass is neither created nor destroyed

• MASS IS CONSERVED IN A REACTION

– Total masses of reactants equals the total mass of products (same is true for atoms on each side of the arrow)

– Remember to include gases too!

• Skeleton equation

– Does not include the relative amounts of the reactants and products involved in the reaction

– First step in balancing a equation

– Shows which elements/compounds are involved

• Ex. Bicycle Equation

– Frame + Wheel + Handlebar + Pedal Bicycle

– To balance this equation, we need to place

quantitative data into our skeletal formula

– 1 Frame + 2 Wheels + 1 Handlebar + 2 Pedals 1 Bicycle

Rules for Balancing Equations

1. Determine the correct formulas for all the reactants and products in the reaction

• May also indicate in parentheses the state in which the reactants and products exist

2. Write the formulas for the reactnats on the left and the products on the right with a yields sign () in between

3. Count the number of atoms of each element in the reactants and products

• A polyatomic ion is counted as a single unit

4. Balance the elements one at a time by using coefficients

• DO NOT CHANGE THE SUBSCRIPTS

5. Check each atom or polyatomic ion to be sure that the equation is balanced

6. Finally, make sure all the coefficients are in the lowest possible ratio

• Law of conservation of energy

– In any chemical or physical process, energy is neither created nor destroyed

– All energy involved can be accounted for as work, stored energy or heat

– Heat (∆H) is transferred from system to surrounding or vice versa

• System – reaction substances

• Surroundings – immediate vicinity around the reaction

– Reference table I

– Endothermic – process that absorbs heat from the surroundings

• System gains heat as the surroundings cool down

• Energy is a reactant (∆H is +)

• Ex. Ice melting, cooking food

– Exothermic – process that releases heat to its surroundings

• System loses heat as the surroundings heat up

• Energy is a product (∆H is -)

• Ex. Hot pack