Chapter 1 Chapter 4
Atoms – The building blocks of matter (H, O, C, Na). A strong electrolyte dissociates completely when dissolved in water. Types of strong electrolytes: solubl
Elements – Made up of atoms of the same kind (O2). ionic compounds, strong acids, strong bases.
Molecules – A particle composed of two or more atoms joined together (CO 2, NaCl). A weak electrolyte only partially dissociates when dissolved in water. Types of weak electrolytes: weak
Compound – Made of two or more different kinds of atoms (CO2, NaCl). acids, weak bases.
All compounds are composed of molecules, but not all molecules are compounds Double Displacement Reactions (double replacement, exchange, metathesis)
Pure Substances – Matter with a fixed composition and unique properties. Elements and compounds are Two ionic compounds exchange ions to form two new compounds. This usually occurs in aqueous soluti
considered pure substances. General Pattern: AX + BY = AY + BX
Mixtures – Comprised of two or more pure substances that exist together but are not chemically If all reactants and products are (aq) then there is NO reaction.
combined. Double displacement reactions will only occur when at least one of the products can be taken out from th
Homogeneous Mixtures: solution.
- Constant, uniform composition throughout. • Forms an insoluble solid (s), precipitate (ppt) (Precipitation Reaction)
- Single phase. o AgNO3 (aq) + NaCl(aq) = AgCl(s) + NaNO3 (aq)
- Examples: air, salt water, metal alloys, gasoline • Forms a liquid (l), for example, water (H2O) (Neutralization Reaction)
Heterogeneous Mixtures: o NaOH(aq) + HCI(aq) = NaCl(aq) + H2O(l)
- Non-uniform composition. • Forms a gas (g), that bubbles out (Gas-forming Reaction)
- May contain multiple phases. o HCI(aq) + NaHCO3 (aq) = CO2 (g) + H2O(l) + NaCl(aq)
- Examples: oil and vinegar, granite, soil, ice water Acid
Physical Properties – Characteristics of a substance that can be observed without changing it into a • Arrhenius: An acid is a substance, when dissolved in water, increases of H+ (or H3O+) ion
different substance (Boiling/melting point, density, mass, color). • Bronsted-Lowry: An acid is a substance that donates a proton to another substance.
Chemical Properties – Characteristics of a substance that can only be observed by changing it into a Base
different substance (Flammability, reactivity, corrosiveness). • Arrhenius: A base is a substance, when dissolved in water, increases the concentration of
Physical Changes – Changes in matter that do not change the composition of a substance but can change OH- ions.
the appearance or physical state. • Bronsted-Lowry: A base is a substance that accepts a proton from another substance.
Chemical Changes – Changes in matter that result in new substances with different properties; also
Loss of elections is oxidation. Gain of elections is reduction. One cannot occur without the other. The
known as chemical reactions. reactions are called redox reaction.
Intensive Properties – Characteristics of a substance that are independent of the amount of substance Oxidation Number - A number assigned to each atom in a compound to represent the “charge” each
present (Boiling/melting point, density, temperature, color, hardness). would have if the electrons were divided among the atoms.
Extensive Properties – Characteristics of a substance that are dependent of the amount of substance Rules for Assigning Oxidation Numbers (ONs):
present (Mass, volume, internal energy, heat change). 1. Pure elements (including those that exist as diatomic) are assigned a zero for their ON.
Chapter 2 N2 (g)
Law of Constant Composition - The composition of a given compound does not change - any sample will N - 0, N - 0
have the same mass percentages of its component elements.
2. Monatomic ions have an ON equal to their common charge.
Law of Conservation of Mass - Mass can be neither gained nor lost, only change arrangement of particles.
• Metals always have a positive charge.
The total mass of substances does not change during a chemical reaction.
• Group 1A metals have an ON = +1.
Law of Multiple Proportions - If elements A and B react to form more than one compound, the different
mass of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers. • Group 2A metals have an ON = +2.
It took about 100 years to discover subatomic particles. Many discoveries led to the fact that the atom itself • Other metals often have a charge equal to their group number on the Periodic
was made up of smaller particles (Elections and cathode rays, radioactivity, nucleus, protons, and neutrons). Table. (Does not apply to transition metals that have more than one oxidation
Atomic Nucleus - Center of the atom; made up of positively charged particles (protons) and uncharged number).
particles (neutrons). 3. The ON of a non-metal depends on what other elements it is bonded to, but can be assigne
Protons - Discovered by Ernest Rutherford in 1919; has a positive charge. Located inside the nucleus. in the following order:
Electrons - Discovered by J.J Thomson in 1897; has a negative charge. Located outside the nucleus. • Fluorine (F) has an ON of -1, ALWAYS.
Neutrons - Discovered by James Chadwick in1932; has no charge. Located inside the nucleus. • Hydrogen (H) has an ON of +1 (Has an ON of -1 when bound to a metal).
Protons and Neutrons make up most of the mass of an atom, Electrons make up the volume. • Oxygen has an ON of -2 (Has an ON of +2 when bound to F and has an ON
Atomic Mass of an element is a “weighted average” of the masses of all isotopes of that element. 1 when in peroxide, and a -1/2 in superoxide).
atomic mass = [(fractional isotope abundance) x (isotope mass)] • Halogens (Group 7A) have an ON of -1.
Example: For naturally occurring carbon 4. The sum of all the ONs in a compound/ion must add up to the overall charge of the
12
C = 12 amu; 98.89% compound/ion. For neutral compounds, this charge is zero.
13
C = 13.0034 amu; 1.11% Molarity
Weighted Average:
[(0.9889) x (12 amu)] + [(0.0111) x (13.0034 amu)] = 12.011 amu
The 7 Diatomic Molecules: Hydrogen (H2), Nitrogen (N2), Oxygen (O2), Fluorine (F2), Chlorine (Cl2), Dilution Equation: Minitial Vinitial = Mfinal Vfinal
Bromine (Br2), and Iodine (I2) Chapter 5
Empirical formulas give the lowest amount whole-number ratio of atoms of each element in a compound Energy
(C6H12O6, H2O2, C6H6). Work (w) - Energy used to cause an object to move against force (w=Fxd)
Molecular formulas give the exact number of atoms of each element in a compound (CH2O, HO, CH). Heat (q) - Energy transferred from a hotter object to a colder object.
When an atom of a group of atoms loses or gains electrons, it becomes an ion. Energy Units
Ion Notations 1 Joule (J) = 1 Newton (N) x meter (m) = 1 J = 1 kg m 2 / s2
• Charges are shown as superscripts after symbol. Internal Energy (E)
• + and - are used to show positive and negative charges.
• For +1 and -1, the one is generally implied.
• Any element written without charge is neutral.
Cations - Form when at least one electron is lost. Monatomic cations are formed by metals. They are
usually written first in a chemical formula. Use the elemental name.
Anions - Form when at least one electron is gained. Monatomic anions are formed by nonmetals, except Enthalpy
the noble gases. They are usually written second in a chemical formula. Use the elemental name with -ide H = E + PV
suffix. The change in enthalpy at constant pressure can be defined by the equation:
Chapter 3 o ΔH = ΔE + PΔV
Guidelines for Balancing Equations Pressure-Volume Work
1. Write correct formulas for reactants and products. • Pressure-volume work (P-V work) is the work involved with the expansion or compressi
2. Balance atoms of each element using coefficients, not subscripts. of gases.
3. Start by focusing on the element that appears in the fewest compounds OR start with the most o w = -PΔV (1 L*atm = 101.3 J)
complex formula compounds. Changes in Enthalpy (ΔH) ΔH = Δ (E + PV)
4. Balance polyatomic ions as a single unit with coefficients if the polyatomic ion(s) exist on • When the system changes at constant pressure, it can be written as:
both sides of the equation. o ΔH = ΔE + PΔV; singe ΔE = q + w and w = -PΔV
Formula Weight (FW) is the sum of the atomic weights for all the atoms in a chemical formula for an o ΔH = (q + w) - w
ionic compound. o ΔH = q
Example: NaCl
• The enthalpy change (ΔH) for a chemical reaction equals the change in internal energy (ΔE
Na: 1 (22.99 amu) + Cl: 1 (35.45 amu) = 58.44 amu
when work (PΔV) is negligible.
Molecular Weight (MW) is the sum of the atomic weights for all the atoms in a chemical formula for a
Enthalpies of Reaction
molecular compound.
• For a chemical change, the change in enthalpy, ΔH, is the enthalpy of the products minus
Example: C2H5OH
enthalpy of the reactants:
C: 2 x (12.011 amu) + H: 6 x (1.008 amu) + O: 1 x (16.00 amu) = 46.07 amu
o ΔHrxn = Hproducts - Hreactants
Percentage Composition from Chemical Formulas
% Element = (number of atoms) x (atomic weight) / FW or MW of the compound) x 100 • ΔHrxn is called the enthalpy of reaction or the heat of reaction.
1 amu = 1.66054 x 10-24 g | 6.02214 x 1023 amu = 1 gram Hess’s Law
Three Methods to Calculate the Limiting Reactant (LR): • If a reaction is carried out in a series of steps, ΔH for the overall reaction equals the sum of
1. Compare moles of reactance you have to moles of reactant required to consume all the other the enthalpy changes for the individual steps.
reactant. • N2(g) + 2O2(g) » 2NO2(g) ΔH = 66.4 kJ
2. Calculate moles of desired product from both (or all) starting materials, whichever produces Standard Enthalpies of Formation: ΔH°f
less product is the LR. • ΔH°f - Standard Enthalpy of Formation
3. Get ratio of moles of each reactant divided by its coefficient from the balanced equation, o Specifically for a formation reaction with reactants and products in their
whichever gives the smallest fraction is the LR. standard states.
• ΔH°rxn - Standard Enthalpy Change
o Obtained from ΔH°f values.
Combination Reaction - When two or more substances react to form one product.
Example: 2 Mg(s) + O2(g) = 2 MgO(s) Calculation: ΔH°rxn from ΔH°f
Decomposition Reaction - When one substance undergoes a reaction to produce two or more other
• So how do we use ΔH°f to calculate unknown ΔH°rxn values?
substances.
Example: CaCO3(s) = CaO(s) + CO2(g)
Combustion Reaction - A rapid reaction that produces a flame.
Example: C3H8(g) + 5 O2(g) = 3 CO2(g) + 4 H2O(g)
Atoms – The building blocks of matter (H, O, C, Na). A strong electrolyte dissociates completely when dissolved in water. Types of strong electrolytes: solubl
Elements – Made up of atoms of the same kind (O2). ionic compounds, strong acids, strong bases.
Molecules – A particle composed of two or more atoms joined together (CO 2, NaCl). A weak electrolyte only partially dissociates when dissolved in water. Types of weak electrolytes: weak
Compound – Made of two or more different kinds of atoms (CO2, NaCl). acids, weak bases.
All compounds are composed of molecules, but not all molecules are compounds Double Displacement Reactions (double replacement, exchange, metathesis)
Pure Substances – Matter with a fixed composition and unique properties. Elements and compounds are Two ionic compounds exchange ions to form two new compounds. This usually occurs in aqueous soluti
considered pure substances. General Pattern: AX + BY = AY + BX
Mixtures – Comprised of two or more pure substances that exist together but are not chemically If all reactants and products are (aq) then there is NO reaction.
combined. Double displacement reactions will only occur when at least one of the products can be taken out from th
Homogeneous Mixtures: solution.
- Constant, uniform composition throughout. • Forms an insoluble solid (s), precipitate (ppt) (Precipitation Reaction)
- Single phase. o AgNO3 (aq) + NaCl(aq) = AgCl(s) + NaNO3 (aq)
- Examples: air, salt water, metal alloys, gasoline • Forms a liquid (l), for example, water (H2O) (Neutralization Reaction)
Heterogeneous Mixtures: o NaOH(aq) + HCI(aq) = NaCl(aq) + H2O(l)
- Non-uniform composition. • Forms a gas (g), that bubbles out (Gas-forming Reaction)
- May contain multiple phases. o HCI(aq) + NaHCO3 (aq) = CO2 (g) + H2O(l) + NaCl(aq)
- Examples: oil and vinegar, granite, soil, ice water Acid
Physical Properties – Characteristics of a substance that can be observed without changing it into a • Arrhenius: An acid is a substance, when dissolved in water, increases of H+ (or H3O+) ion
different substance (Boiling/melting point, density, mass, color). • Bronsted-Lowry: An acid is a substance that donates a proton to another substance.
Chemical Properties – Characteristics of a substance that can only be observed by changing it into a Base
different substance (Flammability, reactivity, corrosiveness). • Arrhenius: A base is a substance, when dissolved in water, increases the concentration of
Physical Changes – Changes in matter that do not change the composition of a substance but can change OH- ions.
the appearance or physical state. • Bronsted-Lowry: A base is a substance that accepts a proton from another substance.
Chemical Changes – Changes in matter that result in new substances with different properties; also
Loss of elections is oxidation. Gain of elections is reduction. One cannot occur without the other. The
known as chemical reactions. reactions are called redox reaction.
Intensive Properties – Characteristics of a substance that are independent of the amount of substance Oxidation Number - A number assigned to each atom in a compound to represent the “charge” each
present (Boiling/melting point, density, temperature, color, hardness). would have if the electrons were divided among the atoms.
Extensive Properties – Characteristics of a substance that are dependent of the amount of substance Rules for Assigning Oxidation Numbers (ONs):
present (Mass, volume, internal energy, heat change). 1. Pure elements (including those that exist as diatomic) are assigned a zero for their ON.
Chapter 2 N2 (g)
Law of Constant Composition - The composition of a given compound does not change - any sample will N - 0, N - 0
have the same mass percentages of its component elements.
2. Monatomic ions have an ON equal to their common charge.
Law of Conservation of Mass - Mass can be neither gained nor lost, only change arrangement of particles.
• Metals always have a positive charge.
The total mass of substances does not change during a chemical reaction.
• Group 1A metals have an ON = +1.
Law of Multiple Proportions - If elements A and B react to form more than one compound, the different
mass of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers. • Group 2A metals have an ON = +2.
It took about 100 years to discover subatomic particles. Many discoveries led to the fact that the atom itself • Other metals often have a charge equal to their group number on the Periodic
was made up of smaller particles (Elections and cathode rays, radioactivity, nucleus, protons, and neutrons). Table. (Does not apply to transition metals that have more than one oxidation
Atomic Nucleus - Center of the atom; made up of positively charged particles (protons) and uncharged number).
particles (neutrons). 3. The ON of a non-metal depends on what other elements it is bonded to, but can be assigne
Protons - Discovered by Ernest Rutherford in 1919; has a positive charge. Located inside the nucleus. in the following order:
Electrons - Discovered by J.J Thomson in 1897; has a negative charge. Located outside the nucleus. • Fluorine (F) has an ON of -1, ALWAYS.
Neutrons - Discovered by James Chadwick in1932; has no charge. Located inside the nucleus. • Hydrogen (H) has an ON of +1 (Has an ON of -1 when bound to a metal).
Protons and Neutrons make up most of the mass of an atom, Electrons make up the volume. • Oxygen has an ON of -2 (Has an ON of +2 when bound to F and has an ON
Atomic Mass of an element is a “weighted average” of the masses of all isotopes of that element. 1 when in peroxide, and a -1/2 in superoxide).
atomic mass = [(fractional isotope abundance) x (isotope mass)] • Halogens (Group 7A) have an ON of -1.
Example: For naturally occurring carbon 4. The sum of all the ONs in a compound/ion must add up to the overall charge of the
12
C = 12 amu; 98.89% compound/ion. For neutral compounds, this charge is zero.
13
C = 13.0034 amu; 1.11% Molarity
Weighted Average:
[(0.9889) x (12 amu)] + [(0.0111) x (13.0034 amu)] = 12.011 amu
The 7 Diatomic Molecules: Hydrogen (H2), Nitrogen (N2), Oxygen (O2), Fluorine (F2), Chlorine (Cl2), Dilution Equation: Minitial Vinitial = Mfinal Vfinal
Bromine (Br2), and Iodine (I2) Chapter 5
Empirical formulas give the lowest amount whole-number ratio of atoms of each element in a compound Energy
(C6H12O6, H2O2, C6H6). Work (w) - Energy used to cause an object to move against force (w=Fxd)
Molecular formulas give the exact number of atoms of each element in a compound (CH2O, HO, CH). Heat (q) - Energy transferred from a hotter object to a colder object.
When an atom of a group of atoms loses or gains electrons, it becomes an ion. Energy Units
Ion Notations 1 Joule (J) = 1 Newton (N) x meter (m) = 1 J = 1 kg m 2 / s2
• Charges are shown as superscripts after symbol. Internal Energy (E)
• + and - are used to show positive and negative charges.
• For +1 and -1, the one is generally implied.
• Any element written without charge is neutral.
Cations - Form when at least one electron is lost. Monatomic cations are formed by metals. They are
usually written first in a chemical formula. Use the elemental name.
Anions - Form when at least one electron is gained. Monatomic anions are formed by nonmetals, except Enthalpy
the noble gases. They are usually written second in a chemical formula. Use the elemental name with -ide H = E + PV
suffix. The change in enthalpy at constant pressure can be defined by the equation:
Chapter 3 o ΔH = ΔE + PΔV
Guidelines for Balancing Equations Pressure-Volume Work
1. Write correct formulas for reactants and products. • Pressure-volume work (P-V work) is the work involved with the expansion or compressi
2. Balance atoms of each element using coefficients, not subscripts. of gases.
3. Start by focusing on the element that appears in the fewest compounds OR start with the most o w = -PΔV (1 L*atm = 101.3 J)
complex formula compounds. Changes in Enthalpy (ΔH) ΔH = Δ (E + PV)
4. Balance polyatomic ions as a single unit with coefficients if the polyatomic ion(s) exist on • When the system changes at constant pressure, it can be written as:
both sides of the equation. o ΔH = ΔE + PΔV; singe ΔE = q + w and w = -PΔV
Formula Weight (FW) is the sum of the atomic weights for all the atoms in a chemical formula for an o ΔH = (q + w) - w
ionic compound. o ΔH = q
Example: NaCl
• The enthalpy change (ΔH) for a chemical reaction equals the change in internal energy (ΔE
Na: 1 (22.99 amu) + Cl: 1 (35.45 amu) = 58.44 amu
when work (PΔV) is negligible.
Molecular Weight (MW) is the sum of the atomic weights for all the atoms in a chemical formula for a
Enthalpies of Reaction
molecular compound.
• For a chemical change, the change in enthalpy, ΔH, is the enthalpy of the products minus
Example: C2H5OH
enthalpy of the reactants:
C: 2 x (12.011 amu) + H: 6 x (1.008 amu) + O: 1 x (16.00 amu) = 46.07 amu
o ΔHrxn = Hproducts - Hreactants
Percentage Composition from Chemical Formulas
% Element = (number of atoms) x (atomic weight) / FW or MW of the compound) x 100 • ΔHrxn is called the enthalpy of reaction or the heat of reaction.
1 amu = 1.66054 x 10-24 g | 6.02214 x 1023 amu = 1 gram Hess’s Law
Three Methods to Calculate the Limiting Reactant (LR): • If a reaction is carried out in a series of steps, ΔH for the overall reaction equals the sum of
1. Compare moles of reactance you have to moles of reactant required to consume all the other the enthalpy changes for the individual steps.
reactant. • N2(g) + 2O2(g) » 2NO2(g) ΔH = 66.4 kJ
2. Calculate moles of desired product from both (or all) starting materials, whichever produces Standard Enthalpies of Formation: ΔH°f
less product is the LR. • ΔH°f - Standard Enthalpy of Formation
3. Get ratio of moles of each reactant divided by its coefficient from the balanced equation, o Specifically for a formation reaction with reactants and products in their
whichever gives the smallest fraction is the LR. standard states.
• ΔH°rxn - Standard Enthalpy Change
o Obtained from ΔH°f values.
Combination Reaction - When two or more substances react to form one product.
Example: 2 Mg(s) + O2(g) = 2 MgO(s) Calculation: ΔH°rxn from ΔH°f
Decomposition Reaction - When one substance undergoes a reaction to produce two or more other
• So how do we use ΔH°f to calculate unknown ΔH°rxn values?
substances.
Example: CaCO3(s) = CaO(s) + CO2(g)
Combustion Reaction - A rapid reaction that produces a flame.
Example: C3H8(g) + 5 O2(g) = 3 CO2(g) + 4 H2O(g)
