MOCK STAAR CHEMISTRY
4A Differentiate between physical and chemical changes and properties


Physical change is change in state (solidLiquidgas)or size or shape
Chemical change: a new substance is formed

Chemical Property: Any characteristic that gives a sample of matter the ability/inability to
undergo a change that alters its composition.
o Examples: Alkali metals react with water; Paper's ability to burn.

Physical property is one that is displayed without any change in composition
4B Identify extensive and intensive properties

Intensive- Properties that do not depend on the amount of the matter present.











Color
Odor
Luster - How shiny a substance is.
Malleability - The ability of a substance to be beaten into thin sheets.
Ductility - The ability of a substance to be drawn into thin wires.
Conductivity - The ability of a substance to allow the flow of energy or electricity.
Hardness - How easily a substance can be scratched.
Melting/Freezing Point - The temperature at which the solid and liquid phases of a
substance are in equilibrium at atmospheric pressure.
Boiling Point - The temperature at which the vapor pressure of a liquid is equal to the
pressure on the liquid (generally atmospheric pressure).
Density - The mass of a substance divided by its volume
Extensive- Properties that do depend on the amount of matter present.




Mass - A measurement of the amount of matter in a object (grams).
Weight - A measurement of the gravitational force of attraction of the earth acting on an
object.
Volume - A measurement of the amount of space a substance occupies.
Length
4C Compare solids, liquids, and gases in terms of compressibility, structure, shape, and volume
4D Classify matter as pure substances or mixtures through investigation of their properties
5A Explain the use of chemical and physical properties in the historical development of the Periodic
Table



Elements in the same column have similar chemical properties
Elements with the same valence electrons exhibit similar properties
Dmitri Mendeleev first discovered the Periodic Table
5B Use the Periodic Table to identify and explain the properties of chemical families, including alkali
metals, alkaline earth metals, halogens, noble gases, and transition metals





Group 1- Alkali metals ( 1 valence electrons, most reactive metals)
Group 2- Alkaline earth metals (2 valence electrons)
Group 17 – Halogens (7 valence electrons, most reactive non-metals)
Group 18 – Noble gases (full valence shell, unreactive or inert)
Transition metals – Groups 3-12
5C Use the Periodic Table to identify and explain periodic trends, including atomic and ionic radii,
electronegativity, and ionization energy
6A Understand the experimental design and conclusions used in the development of modern atomic
theory, including Dalton’s Postulates, Thomson’s discovery of electron properties, Rutherford’s nuclear
atom, and Bohr’s nuclear atom
6B Understand the electromagnetic spectrum and the mathematical relationships between energy,
frequency, and wavelength of light
c = f·λ
As frequency increases, wavelength decreases.
E=h·f (h= Planck’s constant; 6.63 x 10-34 J·s)
As frequency increases, energy increases
6D Use isotopic composition to calculate average atomic mass of an element
 Mass of the isotope (amu) x Abundance on Earth + Mass of another
isotope (amu) x Abundance on Earth

Percentages must be converted to decimal form (68% =.68)
6E Express the arrangement of electrons in atoms through electron configurations and Lewis valence
electron dot structures
7A Name ionic compounds containing main group or transition metals, covalent compounds, acids,
and bases, using International Union of Pure and Applied Chemistry (IUPAC) nomenclature rules
7B Write the chemical formulas of common polyatomic ions, ionic compounds containing main group
or transition metals, covalent compounds, acids, and bases


Cross the “charge” numbers to get subscripts
Aluminum Oxide (Ionic Compound)
7C Construct electron dot formulas to illustrate ionic and covalent bonds
NH3
Step #1: Add up the number of valence electrons that
should be included in the Lewis structure
5 + 3(1) =8
(nitrogen has five; each hydrogen has one)
Step #2: Draw the “skeleton structure” with the
central atoms and the other atoms, each connected
with a single bond
Step #3: Add six more electron dots to each atom
except the central atom. Also, never add dots to
hydrogen
(No Change)
Step #4: Any “leftover” electrons are placed on the
central atom. Find the number of leftovers by taking
the total from Step #1 and subtracting the number of
electrons pictured in Step #3
Step #5: If the central atom has 8, then you are done.
If not, then move two electrons from a different atom
to make a multiple bond. Keep- making multiple
bonds until the central atom has 8 electrons
8-6 = 2 leftover electrons; placed around nitrogen
7D Describe the nature of metallic bonding and apply the theory to explain metallic properties such as
thermal and electrical conductivity, malleability, and ductility

Metals have “loose” valence electrons that leave the individual atoms and forms a “sea of electrons”
7E Predict molecular structure for molecules with linear, trigonal planar, or tetrahedral electron pair
geometries using Valence Shell Electron Pair Repulsion (VSEPR) theory



Two atoms attached to central atom: linear
Three atoms attached to central atom: trigonal planar
Four atoms attached to central atom: tetrahedral
8A Define and use the concept of a mole




The mole is a unit of measure used in chemistry and is used in calculations.
The mole is NOT a measure of mass
1 mole of an atom = atomic mass of the atom
1 mole of a compound = molar mass of that compound
8B Use the mole concept to calculate the number of atoms, ions, or molecules in a sample of material
What is the total number of atoms in 0.260 mol of glucose, C6H12O6?
Given .260 mol
6.02 x 1023
atoms of
C6H12O6
C6H12O6
1.57 x 1023 atoms
of C6H12O6
1 mol
C6H12O6
8C Calculate percent composition and empirical and molecular formulas
To calculate the percent composition of an element in a compound:
1. Find the molar mass of the compound by adding up the masses of each atom
in the compound using the periodic table.
2. Divide the mass of the element by the total molar mass of the compound and multiply by 100.
3. This will give you the percentage for that element.
Steps for Determining an Empirical Formula and Molecular Formula
1. Start with the number of grams of each element, given in the problem.
o If percentages are given, assume that the total mass is 100 grams so that
the mass of each element = the percent given.
2. Convert the mass of each element to moles using the molar mass.
3. Divide each mole value by the smallest number of moles calculated.
4. Round to the nearest whole number. This is the mole ratio of the elements and is represented by
subscripts in the empirical formula.
5. After you have determined the empirical formula, calculate the mass of the compound.
6. Compare that to the molecular mass given.
7. If they are the same empirical formula is the same as the molecular formula
8. If they are different, divide molecular mass by formula mass. Multiply each subscript by that
number.
8D Use the law of conservation of mass to write and balance chemical equations
Coefficients are used to balance equations. You must have the same number of each element in the
reactants and products.
Balance the following
Ca(OH)2 + HCl  CaCl2 + H2O
AgNO3 + CaCl2  Ca(NO3)2 +AgCl
Fe2O3 + C  Fe + CO3
8E Perform stoichiometric calculations, including determination of mass relationships between
reactants and products, calculation of limiting reagents, and percent yield
Na2S2O3(aq) + 4Cl2(g) + 5H2O(aq)  2NaHSO4(aq) + 8HCl(aq)
How many moles of Na2S2O3 are needed to react with 0.12mol of Cl2?
Given: .12 mol ofCl2
1 mol of
Na2S2O3
4 mol of Cl2
.03 mol of
Na2S2O3
In dilute nitric acid, HNO3, copper metal dissolves according to the following equation:
3Cu(s) + 8HNO3(aq)  3Cu(NO3)2(aq) + 2NO(g) + 4H2O(aq)
How many grams of HNO3 are needed to dissolve 11.45g of Cu?
Given: 11.45 g of Cu
1 mol of Cu
63.55 g of Cu
8 mol of HNO3
3 mol of Cu
63.01g of HNO3
1mol of HNO3
30.27 g of HNO3
9A Describe and calculate the relations between volume, pressure, number of moles, and
temperature for an ideal gas as described by Boyle’s law, Charles’ law, Avogadro’s law, Dalton’s
law of partial pressure, and the ideal gas law
P-Pressure (atm, mm Hg, torr)
n- number of moles
V-Volume (L)
T- Temperature (Kelvin)
R- Ideal gas constant (0.0821 L·atm/mol·K; 8.31 L·kPa/mol·K; 62.4 L·mm Hg/mol·K)
Boyle’s law: P1V1 = P2V2
Charles’ Law: V1/T1 =V2/T2
Avogadro’s Law: V1/n1 = V2/n2
Dalton’s Law of partial pressure: Ptotal = P1 + P2 + P3 ….
Ideal Gas Law: PV= nRT
Combined Gas Law: P1V1/T1 = P2V2/T2
9B Perform stoichiometric calculations, including determination of mass and volume
relationships between reactants and products for reactions involving gases
Potassium chlorate is used in the lab to make oxygen gas by the following (unbalanced) reaction:
2KClO3(s) → 2KCl + 3O (g)
(s)
2
How liters of oxygen at STP may be made from reacting 1.226 grams of Potassium chlorate?
Given: 1.23 g of KClO3
1 mol of KClO3
123 g KClO3
3 mol of O2
2 mol of KClO3
22.4 L of O2
1 mol of O2
0.37 L of O2
10B Develop and use general rules regarding solubility through investigations with aqueous
solutions
10C Calculate the concentration of solutions in units of molarity
Molarity = Moles of solute/ Liters of solvent
600.0 mL sample of a 4.50M solution of potassium iodide (KI) contains how many moles of KI?
4.50M = Moles of KI/(.600 L)
4.50(.600) = 2.7 moles ofKI
10E Distinguish between types of solutions such as electrolytes and nonelectrolytes and
unsaturated, saturated, and supersaturated solutions
Unsaturated:
If the solution contains less quantity of solute than what can be dissolved, at a specified temperature it is unsaturated. When more
solute is added into solution the solute dissolves. An unsaturated solution is one in which more of the solute could dissolve at the
same temperature.
Saturated:
A saturated solution is one in which no more of the solute will dissolve at a specific temperature. Solution containing the maximum
amount of solute at a specified temperature is saturated. When more solute is added into the solution the solute will no longer
dissolve.
Supersaturated:
A supersaturated solution is when a solution which contains more solute than would normally dissolve at a certain temperature. The
solution becomes supersaturated by elevating the temperature and the solution contains the maximum amount of solute at an
elevated temperature is carefully cooled to a lower temperature and is now supersaturated. When more solute is added into the
solution, crystals will form.
Electrolytes contain dissociated ions; Nonelectrolytes do not contain dissociated ions
10F Investigate factors that influence solubilities and rates of dissolution such as temperature,
agitation, and surface area
Factors Affecting Solubility:
Temperature affects solubility of solids differently than the solubility of gases.
Solid solubility: ↑ temperature often (not always) ↑ solubility
Solid pressure has no effect on solubility
Gas solubility: ↑ temperature ↓ solubility
Gas solubility: ↑ pressure ↑ solubility
Factors Affecting Rate of Solution Formation (dissolution) of Solids:
↑ temperature ↑rate of dissolution
↑ agitation (stirring or shaking) ↑rate of dissolution
↑surface area (grinding or crushing a solid) ↑rate of dissolution
10G*Define acids and bases and distinguish between Arrhenius and Bronsted-Lowry definitions
and predict products in acid-base reactions that form water
Arrhenius acid-base concept classifies a substance as an acid if it produces hydrogen ions H(+) or
hydronium ions in water. A substance is classified as a base if it produces hydroxide ions OH(-) in water.
Bronsted-Lowry theory classifies a substance as an acid if it acts as a proton donor, and as a base if it acts as
a proton acceptor.
A neutralization reaction which has been put into a word equation: acid + base → salt + water
H2SO4 + 2KOH → 2H2O + K2SO4
H2SO4 + 2KOH → 2H2O + K2SO4
10H*Understand and differentiate among acid-base reactions, precipitation reactions, and
oxidation-reduction reactions
CO2 is the oxidizing agent while CO is reduced
H2 is the reducing agent while H2O is oxidized
10I*Define pH and use the hydrogen or hydroxide ion concentrations to calculate the pH of a
solution
pH = −log[H+]
H+- Hydrogen Ion concentration
10J*Distinguish between degrees of dissociation for strong and weak acids and bases
In general, strong acids and strong bases dissociate, or ionize, completely in aqueous solution. In other words, almost
100 percent of a strong acid or strong base interacts with water and ionizes.
Weak acids and weak bases ionize only slightly in aqueous solution.
Only a very small portion of the molecules will dissociate, representing a small fraction of the entire solution. The same
is true for a weak base; only a small fraction of the molecules will dissociate.
11A*Understand energy and its forms, including kinetic, potential, chemical, and thermal
energies
Kinetic energy- energy provided by molecules that are in motion (Average Kinetic
energy is temperature)
Potential energy- energy that is stored such as the energy stored in chemical bonds
Chemical energy- energy released from chemical compounds
Thermal energy- sum of the kinetic energies in an object
11B*Understand the law of conservation of energy and the processes of heat transfer
11C*Use thermochemical equations to calculate energy changes that occur in chemical reactions
and classify reactions as exothermic or endothermic
Hess’s Law
If change in H is positive, the reaction is endothermic and absorbs heat.
If change in H is negative, the reaction is exothermic and releases heat
11D*Perform calculations involving heat, mass, temperature change, and specific heat
Q = mCpΔT
Q- Energy (Joules)
m- Mass of the substance (kg)
Cp- Specific Heat of the substance (Joules/g°C; Joules/g-K)
ΔT- Change in Temperature (°C or K)
How much heat energy is needed to raise the temperature of a 300 gram piece of copper from 350 º
C to 600 º C? (Specific heat of copper is 0.39 J/g- º C)
Q=?
m= 300. g
Cp= 0.39 J/g- º C
ΔT= Final Temperature –Initial Temperature= 600°C -350°C=250°C
Q = 300(0.39)250 = 29,250 J
12A Describe the characteristics of alpha, beta, and gamma radiation

Alpha particles –helium nucleus with 2 protons and 2 neutrons


Beta particle -- electrons or positrons which do not affect the mass of the element
Gamma radiation – energy with the greatest penetrating power
12C Compare fission and fusion reactions
Fusion: put together
Fission: break apart (chain reaction)
TEKS with * will be assessed for data collection for future tutorials but will not count towards the grade. These are TEKS
that have not been fully covered and assessed through the curriculum scope and sequence.