Chemical Formulas
1. Chemical Symbols: each element has been assigned a one, two, or three letter symbol for its identification. The first letter is ALWAYS CAPITALIZED, any additional letters are always lower case.
2. Diatomic Molecules: When writing the symbols of uncombined elements, almost all are written without a subscript – they are monatomic. Some elements exist in nature as two identical atoms covalently bonded into DIATOMIC MOLECULES. They are: Br2I2N2Cl2H2O2F2 Be sure that whenever you write the formulas for these you write them as diatoms.
3. Chemical Formulas: compounds are composed of combinations of elements chemically combined in definite proportions by weight (mass).
a) Formulas: use chemical symbols and numbers to show both qualitative and quantitative information about a substance.
1) Qualitative: information that cannot be counted or measured (what elements are present)
2) Quantitative: information that can be either counted or measured (how much is present)
b) Types of formulas: Two basic types of formulas provide different types of information about a compound.
1) Empirical Formula: the simplest integer ration in which atoms combine to form a compound CH2
*Formulas for ionic substances are empirical formulas.
2) Molecular Formula: the actual ratio of atoms in a molecule C4H8
4. Atoms, Compounds, and Ions
a) Atoms and compounds are electrically neutral which means that there are equal numbers of positive and negative charges
b) Ions are not neutral and may be either positively or negatively charged. An ion that contains more protons than electrons has an overall positive charge and is called a(n) cation. An ion that contains less protons than electrons has an overall negative charge and is called a(n) anion.
c) Polyatomic Ions (PAI): a group of atoms covalently bonded together possessing a charge. Parentheses are used to enclose PAIwhen there is more than one of the ions in a unit of a compound. The subscript written after the parentheses tells how many of the ions are present in the compound. **The subscript refers to each of the elements in the ion.
Example: (NH4)+ (PO3)-
d) Hydrates: when water from some ionic solutions evaporates, the solute forms a crystal lattice that binds water within the structure. Such a compound is called a hydrate. These crystals have a definite number of water molecules for each unit of the compound. These water molecules are loosely held to the compound. Example: Barium Chloride dihydrate BaCl2 • 2H2O
*The anhydrous (not hydrated) compound can be obtained by heating the crystals to drive off the water. In a chemical reaction, the water in a hydrate does not react. However, it adds mass to the compound.
Energy:
· 2 types of reactions dealing with energy
1. Endothermic:
(inside) (heat)
· Chemical reactions that absorb energy or heat
· It feels cold to the touch
· You’ll see energy added on the left hand side of the equations
i.e. ice packs
2. Exothermic
(outside) (heat)
· Chemical reactions that release energy
· They feel hot to the touch
· You’ll see energy on the right hand side of the equations
i.e. heat packs
Types of reactions
1. Synthesis:
· A combination of 2 or more substances to form a compound
· A + B = C
· i.e. 2 Na (l) + Cl2 (g) = NaCl
2. Decomposition:
· Opposite of synthesis
· AB = A + B
· i.e. 2 H2O (l) = 2 H2 (g) + O2 (g)
3. Single Replacement:
· One element takes the place of another
· A + BC = AB + C
· i.e. Cl2(aq) + 2 KBr (aq) = 2KCl (aq) + Br2 (aq)
4. Double Replacement:
· Elements in compounds change partners
· AB + CD = AC + BD
· i.e.Pb(NO3)2(aq) + 2KI(aq)=PbI2(s)+2KNO3(aq)
5. Combustion:
· Oxygen reacting with a hydrocarbon (CH) to produce Water, Carbon Dioxide and Energy
· CH4 + 2O2 = CO2 + 2H20 + Energy
· Incomplete Combustion is when there is not enough oxygen to react with the hydrocarbon. This produces Water, Carbon Dioxide, Carbon Monoxide, Carbon and Energy
Evidence of a chemical reaction: How do we know a rex has taken place?
· release of a gas (baking a cake)
· color changes (leaves in the fall)
· formation of a precipitate (solid substance formed from solutions)· changes in heat and light (candle burning)
Writing Chemical Formulas
a) Criss-Cross Method: For each element present, write the oxidation state (found in the top right corner of the periodic table). Next, criss-cross those numbers ignoring the sign.
Example: sodium + chlorine
Na+1 + Cl-1 = NaCl
1) sodium + oxygen
Na+1 + O-2 = Na2O
2) calcium + chlorine
Ca+2 + Cl-1 = CaCl2
b) Polyatomic Ions (Table E): Two or more atoms that are chemically combined and possess a net electric charge.
Naming Chemical Formulas
a) Binary Compounds: composed of two elements
I. Metal + Nonmetal: the metallic element is named and written first (it keeps its name), the nonmetallic element is second and its ending is changed to –ide.
II. Nonmetal + Nonmetal: the less electronegative element is named and written first (it keeps its name), the other element is second and its ending is changed to –ide. Prefixes are used to indicate the number of atoms of each nonmetal.
a. Prefixes for nonmetals: 1 mono-
2 di-
3 tri-
4 tetra-
III. Binary Acids: Consist of hydrogen plus some nonmetal.
Prefix hydro- + stem name of the nonmetal + suffix –ic
Example: HCl hydrochloric acid
Example: HBr hydrobromic acid
b) Ternary Compounds: consist of three elements
I. Ternary Bases: Metallic ions combined with the hydroxide ion (OH)
The name of the metallic base followed by the word hydroxide
Example: NaOH is sodium hydroxide
Example: Mg(OH)2 is magnesium hydroxide
II. Ternary Acids: hydrogen ions combined with polyatomic ions
The name of the acid is from the polyatomic ion.
· If the polyatomic ion ends in –ite, the name of the acid ends In –ous.
Example: H2SO4 sulfuric acid
Example: HNO3 nitric acid
· If the polyatomic ion ends in –ate, the name of the acid ends in –ic.
Example H2SO3 sulfurous acid
Example: HNO2 nitrous acid
c) Stock System: In naming compounds of metals that have more than one oxidation number, the name of the metal is followed by a Roman numeral that represents the oxidation number in that compound.
Example: chromium III bromide CrBr3
Example: nitrogen IV oxide NO2
Energy:
· 2 types of reactions dealing with energy
1. Endothermic:
(inside) (heat)
· Chemical reactions that absorb energy or heat
· It feels cold to the touch
· You’ll see energy added on the left hand side of the equations
i.e. ice packs
2. Exothermic
(outside) (heat)
· Chemical reactions that release energy
· They feel hot to the touch
· You’ll see energy on the right hand side of the equations
i.e. heat packs
Types of reactions
1. Synthesis:
· A combination of 2 or more substances to form a compound
· A + B = C *= represents 'yields'
· i.e. 2 Na (l) + Cl2 (g) = NaCl
2. Decomposition:
· Opposite of synthesis
· AB = A + B
· i.e. 2 H2O (l) = 2 H2 (g) + O2 (g)
3. Single Replacement:
· One element takes the place of another
· A + BC = AB + C
· i.e. Cl2(aq) + 2 KBr (aq) = 2KCl (aq) + Br2 (aq)
4. Double Replacement:
· Elements in compounds change partners
· AB + CD = AC + BD
· i.e.Pb(NO3)2(aq) + 2KI(aq)=PbI2(s)+2KNO3(aq)
5. Combustion:
· Oxygen reacting with a hydrocarbon (CH) to produce Water, Carbon Dioxide and Energy
· CH4 + 2O2 = CO2 + 2H20 + Energy
· Incomplete Combustion is when there is not enough oxygen to react with the hydrocarbon. This produces Water, Carbon Dioxide, Carbon Monoxide, Carbon and Energy
Evidence of a chemical reaction: How do we know a rex has taken place?
· release of a gas (baking a cake)
· color changes (leaves in the fall)
· formation of a precipitate (solid substance formed from solutions)
· changes in heat and light (candle burning)
Balancing Chemical Equations notes:
Rules:
1. Check equations by counting atoms on each side of the equation
2. Start by putting 1 in front of the most complex compound
3. Balance elements in multiple reactants or products last
4. Use whole number coefficients (otherwise multiply)
5 Don’t touch the subscripts!!!
Law of conservation of mass:
· matter / energy is neither created nor destroyed
Polyatomics:
· if the same polyatomic appears on both sides of the equation, treat the polyatomic as if it were an atom (a whole unit - don't break it up)
example:
1. 2Li (s) + Br2 (l) = 2 LiBr (s) type of reaction: synthesis
Write Your Questions Below
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What is the mass of 0.0490 mole of potassium sulfate February 25, 2012 |
 | To solve this problem you must follow the steps below:
- find the molecualr mass of 1 mole (K2SO4 K = 2x39 S= 1x32 O =4x16 = 174g/mol
- create a ratio and solve for X 1mol/174g/mol = 0.0490mol/Xg/mol X = 8.526 g
February 25, 2012 - Replied By Expert |
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What is the molar mass of elemental iodine February 25, 2012 |
| Iodine is a diatomic molecule meaning that when not in a compound it will contain two atoms so I2 the mass of I is 127g/mol therefore I2 would equal 254g/mol
Please note the other diatomic molecules Br2I2N2Cl2H2O2F2
February 25, 2012 - Replied By Expert |
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A compound has a molar mass of 118 g/mol and the empirical formula C2H3O2. What is the molecular formula of the compound? February 25, 2012 |
| To solve that problem follow the steps below:
- determine the the empirical mass (in this case it is C = 2x12, H= 3x1, O= 2x16 total=59g/mol)
- divide the molar mass by the empirical mass (in this case it is (118g/mol)/(59g/mol) = 2
- Multiply through the empirical formula with this number so 2(C2H3O2) = C4H6O4
February 25, 2012 - Replied By Expert |
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