MOCK STAAR CHEMISTRY

4A Differentiate between physical and chemical changes and properties

Physical change is change in state (solidLiquidgas)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·λAs 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

8-6 = 2 leftover electrons; placed around nitrogen

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

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 C6H12O6

6.02 x 1023

atoms of 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 Formula1. 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 thatthe 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 formula8. 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

.03 mol of Na2S2O3

4 mol of Cl2

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 8 mol of HNO3 63.01g of HNO3

63.55 g of Cu 3 mol of Cu 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(s) + 3O2(g) 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 3 mol of O2 22.4 L of O2

123 g KClO3 2 mol of KClO3 1 mol 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

0.37 L of O2

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 = mCpΔTQ- 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.