1. Chemistry EOC Review Key Concepts by TEK Must Knows!!! Adapted from Stemscopes, © 2012, Rice University

2. Physical and Chemical Changes and Properties Student Expectation: to differentiate between physical and chemical changes and properties identify extensive and intensive properties classify matter as pure substances or mixtures through investigation of their properties.

3. Physical and Chemical Changes and Properties Key Concepts A physical change to matter will not alter the composition or identity of a substance. A chemical change to matter will always result in the formation of a new substance. The physical properties of matter include properties that describe the substance such as color, smell, boiling point, density or others

4. Physical and Chemical Changes and Properties Extensive properties are dependent on the amount of a substance present, such as mass, number of particles, or energy. Intensive properties are physical properties of matter that are not dependent on the amount of a substance present such as density, ductility, and boiling point.

5. Physical and Chemical Changes and Properties All matter on Earth can be classified as either a pure substance or a mixture. A pure substance consists of a single substance with its own definite composition and properties. A mixture consists of a combination of two or more pure substances with variable composition and properties.

6. States of Matter Student Expectation: compare solids, liquids, and gases in terms of compressibility, structure, shape, and volume.

7. States of Matter Key Concepts Kinetic Molecular Theory states that all matter is composed of particles that are in constant motion. Solid particles have low energy and vibrate in fixed position. Liquid particles have higher energy and are able to move freely past each other. Gas particles are widely spaced with high energy and little attraction to each other

8. States of Matter Solids are relatively incompressible, have a fixed shape, and a definite volume Liquids are more compressible than solids, but much less compressible than gases. Liquids have a definite volume but will take the shape of the container Gases are very compressible and do not have a definite shape. Gases will fill the volume of a closed container. If the container is opened, gases will not have a definite volume.

9. States of Matter PhaseEnergy – move Particle spacing VolumeShapeCompressible SolidLittleClose Rigid Definite No LiquidSomeClose Slide DefiniteNot Definite No GasLotFarNot Definite Not Definite Yes

10. The Periodic Table Student Expectation: to explain the use of chemical and physical properties in the historical development of the Periodic Table 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.

11. The Periodic Table Key Concepts Early scientists, such as Dmitri Mendeleev and Lothar Meyer, observed that the chemical properties of the known elements repeated when placed in order of increasing mass These patterns led to the development of the modern Periodic Table. Group A elements (the representative elements) within the same Group (column) have similar chemical properties.

12. The Periodic Table Group 1A Alkali metals 1 valence electron is lost easily, forms a cation with a 1+ charge. The group has similar physical properties Metallic appearance.

13. The Periodic Table Group 2A Alkaline earth metals 2 valence electrons that can be readily lost, forms a cation with a 2+ charge. The group has similar physical properties Metallic appearance.

14. The Periodic Table Group 7A Halogens (nonmetals) highly reactive with 7 valence electrons, forming an anion with a 1- charge Group 8A noble gases have a complete octet (8) of valence electrons they have little tendency to gain or lose electrons and are non-reactive.

15. The Periodic Table Group B transition metals in the middle of the Periodic Table Number of outer electrons varies, leading to a variety of charges for each element. they share many of their chemical and physical properties The Lanthanide series and the Actinide series, located at the bottom of the periodic table, are the inner transition metals (the rare earth metals).

16. Periodic Trends Student Expectation: to use the Periodic Table to identify and explain periodic trends, including atomic and ionic radii, electronegativity, and ionization energy

17. Periodic Trends Key Concepts Periodic trends patterns that occur across a row (from left to right) or down a column (family) can be used to predict certain properties of elements in their atomic or ionic form.

18. Periodic Trends Atomic radii Decreases when moving from left to right on the Periodic Table due to the increasing number of positive protons within the nuclei pulling on the valence electrons. Increases down a Group due to additional electron shells between the nucleus and repulsion among electrons.

19. Periodic Trends Ionic radii - compared based on their numbers of protons and electrons Across a period, With a greater number of protons, ions will be smaller due to attractive forces between the nucleus and the valence electrons Down a Group, atoms will have an increase in ionic radius for both + and - ions due to the addition of an electron shell. Down a group, a greater number of electrons will cause an ion to become larger due to electron repulsion

20. Periodic Trends electronegativity - ability to attract electrons. increases moving from left to right Decreases moving down a Group, making fluorine the most electronegative element.

21. Periodic Trends Ionization energy - is the amount of energy required to remove an electron from a neutral atom increases moving from left to right Within the same family decreases with increasing atomic number

22. Atomic Theory Student Expectation: to 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.

23. Atomic Theory Key Concepts Concepts of the atom and the nature of matter originated with Greek philosophers more than 2000 years ago. These ideas, though not scientifically tested, formed the basis for later scientists to build on and develop modern atomic theory.

24. Atomic Theory John Dalton (1800’s) investigated the nature of gases in order to gain a better understanding of the laws of conservation of mass and of multiple proportions. His five postulates of atomic theory helped to define the structure and nature of the atom. The scientific community accepted his postulates due to his sound experimental evidence.

25. Atomic Theory J.J. Thomson (late 1800’s) Using cathode rays, he discovered that the rays were actually negatively charged particles with a charge of 1-, and that they were much smaller than atoms. Also studied the relationship between electric charge and matter. Thomson developed the “plum pudding model,” in which electrons were embedded in a positively charged sphere.

26. Atomic Theory Ernest Rutherford (early 1900s) developed the nuclear model of the atom Rutherford’s scattering experiment, (Gold foil experiment) found that atoms contained an extremely small, dense, and positively charged nucleus the area around the nucleus was mostly empty space with a few negative electrons.

27. Atomic Theory Niels Bohr refined the findings of Rutherford Used spectral light emissions to conclude that electrons had specific energy levels His atomic model consisted of spherical shells of electrons on various states surrounding the positively charged nucleus.

28. Atomic Equations Student Expectation: to understand the electromagnetic spectrum and the mathematical relationships between energy, frequency, and the wavelength of light c = λ x f To calculate the wavelength, frequency, and energy of light using Planck's constant and the speed of light. E = hf or E = hc/ λ

29. Atomic Equations Key Concepts The electromagnetic spectrum displays the full range of electromagnetic energy based on wave properties, from high-energy gamma rays to low- energy radio waves. Electromagnetic waves are characterized by energy, frequency, and wavelength Long wavelengths = low frequency and low energy Short wavelengths = high frequency and high energy

30. Atomic Equations c = λ x f All energy waves travel at the same velocity known as the speed of light (c), which equals 3.0 x 10 8 m/s. Wavelength (λ) is defined as the distance between two peaks or two troughs on a wave. Frequency (f) is defined as the number of waves passing a given point per second. Since the speed of light is a constant, frequency and wavelength are inversely proportional

31. Atomic Equations E = hf or E = hc/ λ Planck’s constant equals 6.63 x 10 -34 Js the energy of an electromagnetic wave is directly proportional to its frequency

32. Atomic Structure Student Expectation: to use isotopic composition to calculate average atomic mass of an element express the arrangement of electrons in atoms through electron configurations and Lewis valence electron dot structures.

33. Atomic Structure Key Concepts isotopes - elements with the same number of protons but a different number of neutrons The average atomic mass equals the average of the masses of all the naturally occurring atoms and isotopes for an element.

34. Atomic Structure Electron configuration shows the location and number of electrons in an atom. Energy levels are divided into four sublevels; s, p, d, f. The sublevels are filled with the lowest energy available orbitals filled first. Lewis valence electron dot structures can be used to represent the outer electrons of an atom. Uses the chemical symbol of an element surrounded by dots, each dot represents an electron found in the valence shell.

35. Nuclear Chemistry Student Expectation: to describe the characteristics of alpha, beta, and gamma radiation describe radioactive decay process in terms of balanced nuclear equations compare fission and fusion reaction.

36. Nuclear Chemistry Key concepts Radiation emitted by an element can be characterized as alpha, beta, or gamma radiation Alpha radiation helium nuclei with a +2 positive charge. Beta radiation electron particles with a -1 negative charge gamma radiation high-energy photons with a neutral charge.

37. Nuclear Chemistry Balanced radioactive decay is written to show the conservation of mass number and atomic number during the transmutation of one element into another occurs. Nuclear fission large, unstable atoms split into smaller atoms to achieve a more stable state. Nuclear fusion is the opposite of fission occurs when smaller atoms bind together to form more a larger, more stable atom.

38. Chemical Formulas Student Expectation: to name ionic compounds, covalent compounds, acids, and bases, using International Union of Pure and Applied Chemistry (IUPAC) nomenclature rules write the chemical formulas of common polyatomic ions, ionic compounds, covalent compounds, acids, and bases construct electron dot formulas to illustrate ionic and covalent bonds.

39. Chemical Formulas Key Concepts Ionic compounds name the positive ion (cation) first followed by the negative ion (anion). If the anion is one atom the suffix -ide is added to the anion name. Use a roman numeral for transition metals to indicate the correct charge of the metal cation. The electrons lost by the cation must equal the electrons gained by the anion to form a neutral ionic compound.

40. Chemical Formulas Covalent compounds Use prefixes to show the number of atoms. Then second element ends in -ide. Non-metal atoms share electrons, forming a more stable compound so each atom can achieves a full octet of electrons.

41. Chemical Formulas Acids Binary acids contain just hydrogen and a nonmetal. The prefix hydro- is used followed by the root name of the anion and then suffix -ic. (hydro – ic) Oxyacids (acids that contain hydrogen and a polyatomic ion containing oxygen) ite – ous acid ate – ic acid

42. Chemical Bonding Student Expectation: to construct electron dot formulas to illustrate ionic and covalent bonds describe the nature of metallic bonding apply the theory to explain metallic properties such as thermal and electrical conductivity, malleability and ductility.

43. Chemical Bonding Key Concepts Electron dot formulas are used to show how bonds are formed Ionic bonds, (formula units) Anions and cations are shown in brackets with their respective charges. Lithium oxide

44. Chemical Bonding Covalent bonds, (molecules) the valence electrons are shared between two atoms. Starting with the central atom, the electron pairs are placed around each atom in order to fulfill the octet rule for each atom. Carbon dioxide

45. Chemical Bonding Metallic bonding electrons are delocalized, they do not remain close to any one atom. solid state, the valence electrons of the metal atoms move freely in atom, forming what is known as the “electron sea”. Metallic bonding is formed due to the attraction of the electrons for the metal cations. metallic bonding are more flexible than ionic or covalent bonds metals are ductile and malleable metals are good thermal and electrical conductors.

46. Molecular Structure Student Expectation: to predict molecular structure Linear (180 o ), trigonal planar (120 o ), or tetrahedral (109.5 o ), using Valance Shell Electron Pair Repulsion (VSEPR) theory.

47. Molecular Structure Key Concepts Linear (180 o ), trigonal planar (120 o ), or tetrahedral (109.5 o ) - o r bent

48. The Mole Student Expectation: to define and use the concept of a mole To use the mole concept to calculate the number of atoms, ions, or molecules in a sample of material

49. The Mole Key Concepts The mole is to describe an amount of a substance. The mole is equal to 6.02 × 10 23 atoms, molecules, or formula units of a substance. Is defined by the number of atoms in exactly 12 grams of carbon-12.

50. The Mole One mole = Molar mass The molar mass of a substance (g/mol), can be found by adding the atomic masses of the atoms on the Periodic Table. 1 mole of a gas = 22.4 Liters of a gas (for gases only at STP, standard temperature and pressure)

51. Percent Composition Calculations Student Expectation: to calculate percent composition and empirical and molecular formulas.

52. Percent Composition Calculations Key Concepts Percent composition of a compound represents the percent of each element in a compound by mass. Equals the molar mass of the whole compound divided by the mass of a single element. Then, multiplying by 100 to make a percent. Empirical formula is the lowest whole-number ratio of elements in a compound. Calculated by 1 st : Convert the masses of each element in the compound into mole ratios 2 nd : Divide all by the smallest mole 3 rd : Write the smallest whole number ratios with each element

53. Percent Composition Calculations A molecular formula is the actual chemical formula of a compound Determined based on the molar mass and the empirical formula of the compound. 1 st : The molar mass of the empirical formula is calculated 2 nd : The molar mass of the compound is divided by the molar mass of the empirical formula to find a whole number integer. 3 rd : The empirical formula is then multiplied by this integer to calculate the molecular formula.

54. Chemical Equations Student Expectation: to use the law of conservation of mass to write and balance chemical equations.

55. Chemical Equations Key Concepts Chemical equations must follow the Law of Conservation of Mass - mass is neither created nor destroyed in a non-nuclear change. The total mass of the reactants must equal the total mass of the products. Reactants  products

56. Chemical Equations How to balance 1 st : count the number of atoms of each element on the reactant side and compare that to the number of atoms of the same elements on the product side 2 nd : Use coefficients in front of the chemical formulas to make the number of atoms on each side of the arrow equal.

57. Stoichiometry Student Expectation: to perform stoichiometric calculations determination of mass relationships between reactants and products calculation of limiting reagents percent yield.

58. Stoichiometry Key Concepts Mole ratios are used to determine the relationships among moles in a reaction these ratios come from the coefficients of a balanced chemical equation. The molar masses of reactants and products are used as conversion factors to calculate mass relationships.

59. Stoichiometry A limiting reagent determines the amount of product formed in a reaction, as it is the reactant that is completely consumed first. When one of the reactants is consumed in a chemical reaction, the reaction stops and no further products may be formed. Use stoichiometry to calculate the limiting reactant determine which reactant is consumed first, and then use the amount of this reactant to find the moles or mass of product formed.

60. Stoichiometry Percent yield of a reaction is found by dividing the actual amount of a product by the theoretical amount of the product, and then multiplying by 100 to make a percent. % yield = actualx 100 theoretical Stoichiometry is used to find the theoretical (ideal) yield of a chemical The actual yield is a measurement of the actual amount of product made during a reaction.

61. Gas Laws Student Expectation: to describe and calculate the relations between volume, pressure, number of moles, and temperature for an ideal gas as described by Boyles's law, Charles' law, Avogadro's law, Dalton's law of partial pressure, and the ideal gas law describe the postulates of kinetic molecular theory.

62. Gas Laws Key Concepts Kinetic Molecular theory is used to describe the behavior of ideal gases. ideal gas particles are of negligible size compared to the space between them the particles are in continuous, rapid, random motion their collisions are elastic there are no significant interactions among the particles of a gas When temperature increases for an ideal gas, the kinetic energy of the particles increases proportionally.

63. Gas Laws Boyle’s law, P 1 V 1 = P 2 V 2 the pressure exerted by gas particles is inversely proportional to the volume occupied by the gas. At constant temperature, an increase in pressure will result in a decrease in volume.

64. Gas Laws Charles’s law, V 1 /T 1 = V 2 /T 2 The temperature of a gas is directly proportional to the volume occupied by the gas. When the pressure is held constant, a decrease in temperature results in a decrease in volume. temperature must be in Kelvin.

65. Gas Laws Avogadro’s law, V 1 /n 1 = V 2 /n 2 relationship between the number of moles (n) of a gas, and the volume occupied by the gas at a constant temperature and pressure, where the volume of a gas is directly proportional to the number of moles of a gas. Avogadro’s principle states that equal volumes of gases contain the same number of atoms or particles.

66. Gas Laws Dalton’s law of partial pressures, P total = P 1 + P 2 + P 3 +…P n The total pressure exerted by a mixture of gases is equal to the sum of the individual pressures of all the gases in the mixture.

67. Gas Laws The ideal gas law, PV = nRT describes the relationship between the temperature, pressure, volume, and the number of moles of a gas under specific conditions. This law represents a combination of the relationships described in Boyle’s law, Charles’ law, and Avogadro’s law. The ideal gas constant R is a physical constant and is specific to the units used for the pressure and volume. Temperature must be in Kelvin.

68. Gas Stoichiometry Student Expectation: to perform stoichiometric calculations including determination of mass and volume relationships between reactants and products for reactions involving gases.

69. Gas Stoichiometry Key Concepts Using a balanced chemical equation, stoichiometric ratios, and molar masses to determine the mole and mass relationships between the amounts of products and reactants in a chemical equation involving gases. For reactions not at STP, stoichiometric ratios and the Ideal Gas Law can be used to find mole and volume relationships between products and reactants.

70. Aqueous Solutions Student Expectation: to describe the unique role of water in chemical and biological systems develop and use general rules regarding solubility through investigations with aqueous solutions.

71. Aqueous Solutions Key Concepts Water has several unique properties. amphoteric, (can be an acid or a base) is highly polar, has a higher boiling and melting point than other compounds of the same structure due to the hydrogen bonding between water molecules. The strong attraction between water molecules (hydrogen bonding) accounts for properties such as cohesion and surface tension.

72. Aqueous Solutions Life on Earth is highly dependent on the unique properties of water most of the metabolic processes in biological organisms take place in aqueous solutions. water is an essential component of the energy transformation processes of photosynthesis and cellular respiration. water’s high specific heat plays a critical role in Earth’s relatively moderate temperature variations.

73. Use STAAR Chart to determine Which Ionic compounds are soluble (aq) and insoluble (s)

74. Molarity Student Expectation: to calculate the concentration of solutions in units of molarity use molarity to calculate the dilutions of solutions

75. Molarity Key Concepts Molarity (M) is used to express the concentration of a solution moles of a solute dissolved per liter of solution, or mol/L. This unit of mol/L is also called “molar.” Diluted solutions may be created from concentrated solutions by merely adding more solvent. M 1 V 1 = M 2 V 2,

76. Solutions and Solubility Student Expectation: to distinguish between types of solutions such as electrolytes and nonelectrolytes and unsaturated, saturated, and supersaturated solutions investigate factors that influence solubilities and rates of dissolution such as temperature, agitation, and surface area.

77. Solutions and Solubility Key Concepts When ionic compounds dissolve in a solvent, the charged ions in the solution can conduct an electric current and are called electrolytes. When molecular compounds dissolve in a solvent, the molecules in solution do not conduct electric current and are called nonelectrolytes.

78. Solutions and Solubility saturated solution - a solution in which the maximum amount of solute is dissolved in the solvent supersaturated - the solution contains more than the maximum amount of solute unsaturated - it contains less than the maximum amount.

79. Solutions and Solubility Factors that affect the solubility of a solute in a solvent temperature of the solvent intermolecular forces (for gases) the partial pressure of a gas over a liquid solute. Rates of dissolution of a solute in a solvent depend on Temperature Agitation (stirring) surface area of the solute

80. Acids and Bases Student Expectation: to define acids and bases and distinguish between Arrhenius and Brønsted-Lowry definitions Understand, differentiate, and predict products in acid base reactions, precipitation reactions, and oxidation-reduction reactions define pH and use the hydrogen or hydroxide ions concentrations to calculate the pH of a solution distinguish between degrees of dissociation for strong and weak acids and bases.

81. Acids and Bases Key Concepts The Arrhenius definition of acids and bases in aqueous solutions, an acid forms more hydrogen ions (H + ) a base forms more hydroxide ions (OH - ) The Brønsted-Lowry definition of acids and bases an acid can donate a proton (H + ) a base can accept a proton.

82. Acids and Bases Acid base reactions involve the transfer of protons. In these neutralization reactions, the acid and the base normally react to form water and salt. A precipitation reaction occurs when an insoluble precipitate is formed. Both precipitation and acid-base reactions are types of double replacement reactions. Oxidation-reduction (redox) reactions involve the transfer of electrons

83. Acids and Bases The pH scale is used to determine the strength of an acid or base pH = – log [H + ] The ranges on the pH scale below 7 for acids above 7 for bases = 7 for neutral solutions [H + ] x [OH - ] = 1 x 10 -14

84. Acids and Bases Complete dissociation of ions occurs with strong acids and strong bases, partial dissociation of ions occurs with weak acids and weak bases. The conjugate base of a strong acid is a weak base. The conjugate acid of a strong a base is a weak acid.

85. Energy Student Expectation: to understand energy and its forms, including kinetic, potential, chemical, and thermal energies understand the law of conservation of energy the processes of heat transfer.

86. Energy Key Concepts Matter may contain one or more types of energy, where energy is defined as the ability to do work. Kinetic energy (energy of motion) based on the movement of an object or a substance. Potential energy (energy of position) stored energy based on the composition of a substance the position of an object in space.

87. Energy Chemical energy (form of potential energy) the energy stored in the bonds between the atoms of molecules and ionic crystals. the potential of a chemical substance to undergo a transformation through a chemical reaction the ability of a system (the chemical reaction) to do work during chemical reactions. During chemical reactions, energy is transferred, either producing thermal energy or requiring thermal energy.

88. Energy Thermal energy is the total internal energy that a substance possesses as well as the total KE of the particles in a system. It depends on the temperature of that system relates to both heat and temperature. Temperature is a measure of the average KE within the system heat is the movement of thermal energy from one substance to another.

89. Energy The law of conservation of energy states that energy can neither be created nor destroyed, but instead is transferred between a system and its surroundings. This energy transfer often occurs in the form of heat flow. Heat, or the movement of thermal energy from a warmer substance to a cooler substance, flows between a system and its surroundings

90. Thermochemical Equations Student Expectation: to use thermochemical equations to calculate energy changes that occur in chemical reactions classify reactions as exothermic or endothermic perform calculations involving heat, mass, temperature change, and specific heat

91. Thermochemical Equations Key Concepts A thermochemical equation is a balanced chemical equation that includes the amount of energy absorbed or released during a chemical reaction. The total heat energy of the system is known as the enthalpy (H) of the system. ΔH is often expressed in kilojoules (kJ).

92. Thermochemical Equations Exothermic reaction, ΔH is negative heat flows from the system to the surroundings. The energy released during the reaction is represented as a product in the thermochemical equation. Endothermic reaction, ΔH is positive heat flows into the system from the surroundings. The energy absorbed during the reaction is represented as a reactant in the equation.

93. Thermochemical Equations The ΔH (change in enthalpy) of a system is equivalent to the heat of the system (q) q = mc ΔT q = m (mass in grams) x c (specific heat) x ΔT (change in temperature). The specific heat (c) of a substance is defined as the amount of heat required to raise the temperature of 1 gram of the substance by 1 degree Celsius.

94. Calorimetry Student Expectation: to use calorimetry to calculate the heat of a chemical process.

95. Calorimetry Key Concepts Calorimetry includes the complete set of equations and experimental procedures used to measure the heat flow for physical and chemical processes. A calorimeter is the device used to measure the heat absorbed or released during a chemical process.