Transcript UNIT 4: Formulas and Equations (Review Book Topic 2)

UNIT 4: Formulas and
Equations
(Review Book Topic 2)
How can we distinguish between quantitative and qualitative
information?
What are the different types of formulas?
How can we know that elements form a compound and in
what proportions?
How can we write formulas?
How can we name compounds?
What are the different parts of a chemical equation?
What are the differences between endothermic and exothermic
reactions?
How can chemical equations demonstrate the Law of
Conservation?
What are the different type of chemical reactions?
How can we predict products of a chemical reactions?
How can we determine an unknown reactant, product, or mass
in a chemical equations?
AIM: How can we distinguish
between quantitative and
qualitative information?
 Chemical Symbols: Each element has a unique one-,
two-, or three- letter symbol
 The first letter is always capitalized (Table S, PT)
 Almost all symbols are written without a subscripts as
monatomic
 Diatomic Molecules: elements that exist in nature as
two identical atoms covalently bonded into a diatomic
molecule
 Br2, I2, N2, Cl2, H2, O2, F2
AIM: How can we distinguish
between quantitative and
qualitative information?
 Chemical Formulas: compounds are composed of elements
chemically combined in fixed ratios

Formulas use chemical symbols and number to show both
qualitative and quantitative information about a substance
 Qualitative: information that relates to things that cannot be
counted or measured
What elements are in the compound
 Quantitative: information that deals with things that can
either be counted or measured

The number of atoms of each element in the compound
(subscript or coefficient can give us this info)
AIM: How can we distinguish
between quantitative and
qualitative information?
Example: Determine the quantitative and qualitative
information in the following examples:
1.
CaCO3
2.
Zn3(PO4)2
AIM: What are the different types
of formulas
 Empirical: simplest ratio
 Ionic formulas are always empirical formulas
 Molecular: actual ratio
 covalently bonded substances form molecules, in some
cases the empirical represents both empirical and
molecular –H2O-
 In other cases molecular formula is a multiple of the
empirical formula C6H12O6 – six times the empirical
CH2O
AIM: How can we know what
elements form a compound and in
what proportions?
 Atoms and compound are electrically neutral – equal
numbers of negative (electrons) and positive (protons)
 Ions – can be positive or negative
 Ionic charge: indicated by a superscript positive – lost
electrons , negative – gained electrons Cl-, Al+3
 Polyatomic Ion: group of atoms covalently bonded
together, possessing a charge – Table E
AIM: How can we know what
elements form a compound and in
what proportions?
 Forming a compound: many ways – one way:
 By attraction of oppositely charged ions
 Monatomic or polyatomic ions attract each other in a
ratio that produce a neutral compound
 Coefficients: written in front of a formula, applies to
entire formula, multiple coefficient and subscript to
determine the number of each type of element
AIM: How can we know what
elements form a compound and in
what proportions?
 Hydrates
 Compounds that contain definite
amounts of water molecules
 Ex: BaCl2·2H2O – Barium
chloride traps 2 water molecules
AIM: How can we write
formulas?
 Compounds must be electrically neutral
 For many elements, the oxidation state is equal to the charge
found in the top right corner of each element box
 1:1
 Na+ & Cl - yields NaCl
 Mg2+ & S2- yields MgS
 Not 1:1 – Mg2+ & Cl
-
 Write the charge of one ion as the subscript of the other
without the sign (# only)
 Thus MgCl2; that is 1 Mg with a 2+ & 2 Cl with 1- each yields
(2+) + 2(1-) = 0
AIM: How can we write
formulas?
Examples: Write the formulas for the following:
a. Al+3 and Br –
b. Ba+2 and CO3 -2
c. Cu 2+ and CO3 -2
d. Pb +2 and Cl –
e. Pb +4 and CrO4 -2
AIM: How can we name
compounds?
 Binary Ionic (Metal–positive & Nonmetal–negative)
 1st element (metal) – retains its name
 2nd element (nonmetal) – change ending to “ide”
 Ex: KCl is Potassium chloride
 Other Ionic (Contains Polyatomic Ions)
 Same as Binary except all polyatomic ions retain their name
 Examples:
Potassium nitrate
 KNO3 is __________________
Ammonium chloride
 NH4Cl is _____________________
ammonium nitrate
 NH4NO3 is___________________
AIM: How can we name
compounds?
Examples: Write the names for the following formulas:
a. Mg(SO4)
b. Na(OH)
c. Ca(OH)2
d. Li3(PO4)
e. (NH4)Cl

AIM: How can we name
compounds?
Covalent (2 Nonmetals)
 Need prefixes to tell reader how many of each
 Exception – if only one 1st element, don’t use mono
 Examples
Nitrogen monoxide
 NO is _____________________
Dinitrogen tetroxide
 N2O4 is _____________________
 Stock System (multiple oxidation states)
 mostly the touchable metals; Iron, Tin, Copper . . .
 a Roman numeral proceeding the metal tells the reader the
oxidation number
 Examples
Iron (II) chloride
 FeCl2 is _________________
Iron (III) chloride
 FeCl3 is _________________
AIM: What are the different
parts of a chemical equation?
 Reactant(s) yield Product(s)
 Reactant are on the left of the arrow (yield sign) and
products are to the right of the arrow
 State of matter is indicated by the letter inside the
parenthesis
 Example: C (s) + O2 (g) → CO2 (g)
 Identify the reactants and products in the equation
above…….
AIM: What are the differences between
endothermic and exothermic reactions?
 Endothermic
 Heat is required for a reaction to occur, thus, energy
is found on the reactant side
 Example: H2O (s) + energy → H2O (l)
 Exothermic
 Heat is produced in a reaction, thus energy is found
on the product side
 Example: H2O (g) → H2O (l) + energy
AIM: How can chemical equations
demonstrate the Law of
Conservation? - BALANCING
 The Law of Conservation of Mass & Charge must
be upheld
 Example:
1 O2 (g) → __
2 H2 (g) + __
 __
2 H2O (g)
 Remember – It’s coefficient x subscript to find the
# of atoms
AIM: How can chemical equations
demonstrate the Law of
Conservation? - BALANCING
 Nothing can be created or destroyed – law of conservation of
mass
 Count up the atoms on both side and fill in any missing elements
or compounds
 Same thing for missing mass
 Practice:
If 103.0g of potassium chlorate are decomposed to form 62.7g
of potassium chloride and oxygen gas according to the
equation
2KClO3  2KCl + 3O2 how many grams of oxygen are formed?
AIM: How can we determine an
unknown reactant, product, or mass in a
chemical equations?
 Nothing can be created or destroyed – law of
conservation of mass
 Count up the atoms on both side and fill in any missing
elements or compounds
 Same thing for missing mass
AIM: What are the different
types of chemical reactions?




Synthesis

2 or more reactants form 1 product

A + B → AB
Decomposition

1 reactant breaks down into 2 or more products

AB → A + B
Single Replacement

1 element replaces another

A (element) + BX (compound) → B (element) + AX (compound)
Double Replacement

2 elements/polyatomic ions replace two others

AB (compound) + CD (compound) → AD (compound) + CB (compound)
AIM: How can we predict
products of a chemical reactions?
 Single Replacement Reactions:
 If the individual metal is above the metal that is in
the compound a reaction will occur
 Double Replacement Reactions:
 If a solid is formed (Table F)
 If a gas is formed
 If a molecular substance such as water is formed
Using TABLE F
1.
Cross out the first element or compound
2.
Look for second element or compound on Table F check to
see if it is paired with an exception
3.
Determine solubility
Ex)
NH4Cl 
MgOH
Soluble
Insoluble
TABLE F
AgNO3(aq) + NaCl(aq)  AgCl(_____)
+ NaNO3(____)
aq
s