Types of Measurement

1. Types of Measurement Qualitative measurement: uses words ex. A yellow-green gas was released. Quantitative measurement: uses numbers ex. The oxide has a mass of 1.567 grams.

2. TEMPERATURE SCALES Celsius Kelvin Fahrenheit Compare all the scales Temperature Conversions

3. Temperature Scales Converting Celsius Temperature to Kelvin Temperature: C + 273 = ____K Example: 25 C = ______ 25 + 273 = 298 K Converting Kelvin to Celsius: Reverse K - 273 = ______C 300 K = _________C 300 � 273 = 27 C

4. Law of Conservation of Mass Discovered by Antoine Lavoisier Mass is neither created nor destroyed Combustion involves oxygen, not phlogiston

5. Classification of Matter

6. Matter:  Anything occupying space and having mass.

7. Element: A substance that cannot be decomposed into simpler substances by chemical means.

8. Types of Mixtures Mixtures have A homogeneous variable composition. mixture is a solution (for example, vinegar) A heterogeneous mixture is, to the naked eye, clearly not uniform (for example, a bottle of ranch dressing)

9. Mixtures Can be isolated by separation methods: ? Chromatography ? Filtration ? Distillation

10. a) compound b) element c) homogeneous mixture d) heterogeneous mixture 1. Concrete 2. air 3. salt 4. gold 5. helium 6. tea 7. sea water

11. Metric Prefixes Powers of 10 Know the Metric Prefix to Power of 10 Mega (106) to micro (10-6) 6000mg ______g______cg_____kg ________Mg

12. Precision vs. Accuracy PRECISION: Replication of results ex. A penny is massed 3 times ---- 2.6g, 2.6g, 2.6g ACCURACY: True or correct results

13. Seven Base SI Units Length meter (m) Mass kilogram (kg) Time second (s) Temperature Kelvin (K) Current ampere (amp) Amount mole (mol) Luminous Intensity candela (cd)

14. DERIVED UNITS A combination of two or more units. Examples: Speed miles/hr ft/sec DENSITY is the mass/volume. Area = length x width

15. Density Density = Mass/Volume (g/mL or g/cm3) Water is the standard for all density values = 1.0 g/mL V = M/D (1mL=1cm3) M= D x V (g) What is the volume of a 50 g metal block with a density of 5 g/cm3 ?

16. Density Density is the mass of substance per unit volume of the substance:

17. Mass Mass = Density x Volume

18. Volume Volume = Mass / Density

20. The Chemists� Shorthand: Atomic Symbols

21. Atomic Number = # of Protons Also represents # of electrons Mass Number = Total # of Protons and neutrons Sometimes the mass number and atomic mass are interchanged

22. Neutrons and Protons are located in the nucleus

23. Atom is mostly empty space

25. The Mass and Change of the Electron, Proton, and Neutron

26. Periodic Table Groups ??properties ? atomic number Groups (vertical) 1A = alkali metals 2A = alkaline earth metals 7A = halogens 8A = noble gases Periods (horizontal) Periods 1-7

28. Organization of the Periodic Table A. Metals vs Nonmetals - see staircase on the periodic table 1. Metals - are to the left of the staircase; most are solids; conduct electricity; lose electrons and form positive ions

31. Metals vs Nonmetals 2. Nonmetals- are to the right of the staircase; most are gases; nonconductors of electricity; gain electrons and form negative ions

33. Metals vs Nonmetals 3. Metalloids (Semimetals) - in purple in backside of textbook. These elements border the staircase and have properties of both metals and nonmetals.

34. Periods There are 7 main periods on the periodic table. They are numbered # 1-7. They represent the major energy levels (n). The periods are horizontal rows that extend from left to right. Ex: Period 2 includes Li - Ne.

36. Groups or Families 1. Groups- vertical row of elements . A. IA - called the Alkali Metals (1 valence electron) very reactive B. IIA - called the Alkaline Earth Metals (2 valence electrons)

37. Groups or Families 1. Groups- vertical row of elements . A. IA - called the Alkali Metals (1 valence electron) very reactive B. IIA - called the Alkaline Earth Metals (2 valence electrons)

39. Groups or Families C. VIIA - Halogen group (has 7 valence electrons) very reactive D. VIIIA or Group O - Noble, Rare, or Inert Gases (has 8 valence electrons except for Helium) nonreactive-very stable

41. Representative vs Transition Elements 1. Representative Elements - The Group �A� Elements which include all the Groups IA to VIIIA 2. Transition Elements- The Group �B� Elements

43. Inner Transition Elements A. Lanthanide Series - the 4f row that includes # 57 (Lanthanum) through #71 Lu B. Actinide Series - the 5f row that includes #89 Ac (Actinum) through #102 No

45. Periodic Trends 1. Atomic Radius 2. Electronegativity 3. Ionization Energy

46. Atomic Radius Atomic Radius is defined as the distance from the center of the nucleus to the outermost valence shell Periodic Trend: Down a group the atomic radius increases. Across a period it decreases Why?

47. Atomic Radius

48. Common Ions Ion: a positively charged atom Cation (+) ion Anion (-) ion

49. Figure 2.22 Common Cations and Anions

50. Classification of Inorganic Compounds There are 2 main kinds of compounds 1. Ionic: made up of ions of opposite charge . The strong electrostatic force of attraction between them is called the ionic bond. Electrons are TRANSFERRED 2. Covalent : made up of 2 or more nonmetals that SHARE pairs of electrons between their nuclei.

51. Chemical Bonds Why do atoms bond? Atoms seek to become chemically stable. To do this their valence shell must be complete. The Octet Rule states that atoms will either gain, lose , or share valence electrons to attain �8� electrons in their valence shell to become stable. There is only one group of elements that are already stable and that is the Noble Gases

52. Ionic Compounds Ions of opposite charge 1. Metal cation (+) is written first and is named by the metal�s name 2. Nonmetal anion (-) is written second and is named by the nonmetal�s name with a revised ending of -ide.

55. Metal (+) and Nonmetal (-) Ions

56. Examples of Binary Ionic Compounds 1. Sodium chloride 2. Lithium nitride 3. Barium phosphide

57. Naming Compounds 1. Cation first, then anion 2. Monatomic cation = name of the element Ca2+ = calcium ion 3. Monatomic anion = root + -ide Cl? = chloride CaCl2 = calcium chloride

58. Binary Ionic Compounds Writing and naming these compounds 1. Cation (+) first; Anion (-) second 2. Net charge of ions = 0. 3. Subscripts used to indicate the # of ions needed to attain net charge = 0.

59. Examples of Binary Ionic compounds 1. Al2S3 2. BaO 3. MgBr2 Name these compounds

61. Covalent or Molecular Compounds Made up of nonmetals that �SHARE� electrons between atoms. This type of bond is called a covalent bond.

62. NAMING BINARY COVALENT 1. First nonmetal�s name is that of the elements. 2. Second nonmetal�s name has an -ide ending (just like ionic) 3. Use prefixes to describe the subscripts (1-mono; 2-di; 3-tri; 4-tetra; 5-penta; 6-hexa; 7- hepta; 8-octa; 9-nonea; 10- deca)

63. Covalent or Molecular P2O5 = dphosphorus pentoxide How would you write Carbon monoxide Tetranitrogen decoxide

64. Balancing Chemical Reactions The Law of Conservation of Matter states that the mass of the products is equal to the mass of the reactants (Matter is not created or destroyed) Balancing chemical equations is an abbreviated form to insure that a chemical reaction obeys the Law of Conservation of Matter.

65. 5 Guidelines to Balance Equations 1. Count the number of each element in the reactant and product side. 2. Use COEFFICIENTS (numbers in front of the chemical symbol or formula) 3. Never add or change subscripts. 4. There are 7 Diatomic elements (N2,O2,F2,Cl2,Br2,I2,,H2) 5. Balance Hydrogens and Oxygens last

66. BALANCING CHEMICAL EQUATIONS Symbolic language used to describe a chemical reaction Equation means �EQUAL� The Law of Conservation of Matter states the mass of the products MUST be equal to the mass of the reactants Quantities of Reactants and Products are expressed in moles by using Coefficients

67. SYMBOLS IN REACTIONS (s) - solid (g) - gas (l) - liquid (aq) - aqueous�.. dissolved in water See board for other symbols

68. Practice Balancing Equations Go to www.usaprep.com and practice balancing equations Try this one: H2 + O2 yields H2O

69. Balanced Equation

70. Acids and Bases Operational Definitions of Acids: 1. Tastes sour 2. Neutralizes the actions of bases 3. Blue litmus turns red 4. Liberates Hydrogen gas when reacted with certain metals Examples: Foods and drinks

71. Acids and Bases Operational Definition of Bases 1. Tastes bitter 2. Slippery to touch 3. Red litmus turns blue 4. Neutralizes the action of acids Examples: Cleaning solutions

72. Neutralization Reaction Acid + Base yields salt + water Example: HCl + NaOH yields NaCl + H2O

73. Strength, Corrosiveness, and Concentration STRENGTH of an acid or base is defined by how much it ionizes in solution HCl (Hydorchloric Acid) ionizes almost 100% in solution so it is considered a very STRONG acid HC2H3O2 (Acetic Acid) ionizes only 1 % so it is considered a very WEAK acid

74. STRONG ACIDS Hydrochloric Sulfuric Nitric

75. Weak Acid and a Weak Base Base: Ammonia Acid: Acetic Acid

76. pH Scale The pH scale has values from 0 - 14. 0-6 is considered Acidic 7 is neutral 8-14 is basic

77. The Mole A mole is the SI unit of measure for the amount of a substance. Chemists needed a way to measure the mass of elements and compounds .

78. 3 Ways to Measure and Define the Mole I. Mass II. Number III. Volume

79. Gram Atomic Mass The atomic mass of an element expressed in grams One mole of an element is its gram atomic mass

80. The mass of one mole of K (Potassium) is____________________

81. The Chemists� Shorthand: Atomic Symbols

82. The mass of one mole of K (Potassium) Is 40 grams

83. PHYSICAL AND CHEMICALCHANGES 1. Physical Change- only the appearance changes ; the identity is still the same. What changes? A. Size B. Shape C. Phase or State

84. EXAMPLES 1. Physical changes: a. Size or shape : splitting, dissolving, breaking, or tearing. b. Phase Change : melting, vaporizing, freezing, or condensing.

85. EXAMPLES 2. Chemical Changes : rusting, growing, burning, combusting, produces, reacts, fermenting, cooking, frying, and exploding. (Any action that would result in a new product )

86. 2. Chemical change : The appearance and identity changes. A new product is formed. How do you know if a new product is formed ? A. Gas is formed B. Color change C. Change in mass D. Heat change E. Solid formed F. Light released

87. How do you know if a new product is formed ? A. Gas is formed B. Color change C. Change in mass D. Heat change E. Solid formed F. Light released

89. NUCLEAR REACTIONS An unstable nucleus breaks down and emits radioactive particles. 3 Types of Radioactive Decay

90. Alpha Decay An alpha particle is a Positively charged particle. Is actually a Helium atom that has lost 2 electrons. Has a (+) 2 charge Largest, slowest, and less penetrating particle

91. Beta Decay A negatively charged particle Mass of an electron Is basically a fast accelerated electron More penetrating than an alpha particle

92. Gamma Radiation Has no mass or charge Is not a particle Is a form of energy Very penetrating (can be shielded by lead) Most damaging

93. HALF LIFE The half life of a radioisotope is the time it takes for one half of that isotope to decay.

94. Example: The half-life of Mercury-195 is 31 hours. If you start with 20 g, how much will be left after A) 31 hours? B) 62 hours?

95. A) 10 grams B) 5 grams

96. Particles in an Electric Field

97. FUSION and FISSION Fission is the splitting of nuclei resulting in a tremendous release of energy Fusion is the combining of nuclei resulting in even more energy released. The sun produces energy as a result of nuclear fusion ( 2 H atoms combine to form Helium)

98. Unit 10 Energy/ Heat/ Phase Changes Temperature is the measure of the average kinetic energy of particles. Heat is a form of Energy that may be absorbed or released. Heat flows from a warm body to a cooler one until equilibrium is reached. A calorie is defined as the amount of heat required to raise the temperature of 1 g of water one degree Celsius. (unit)

99. Types of Phase Changes Melting: solid to liquid Freezing: liquid to solid Evaporation- liquid to gas Condensation- gas to liquid Sublimation - solid to gas Deposition- gas to solid

100. Key Terms for Heat 1. Energy - the capacity to do work 2. Heat - Energy that is transferred from one object to another 3. Thermochemistry - study of heat effects in chemical reactions 4. Combustion - reactions that release heat

101. Key Terms 5. Exothermic reaction - reactions that release heat 6. Endothermic reactions - reactions that absorb heat

102. Heat The SI unit of Heat is a Joule 4.18 joules = 1 calorie

103. Specific heat capacity is the amount of heat required to raise the temperature of one gram of a substance one degree Celsius. The specific heat capacity of water is 1 . Water has a high specific heat capacity due to hydrogen bonds Metals have low specific heat capacities

104. Energy in Chemical Reactions Bond breaking in reactants requires energy. Bond formation for new products releases energy. An endothermic or exothermic reaction is determined by the balance between these two processes.

105. Phase Changes There are 2 types of energy utilized in a phase change problem Potential Energy is used DURING a Phase Change Kinetic Energy is used when a temperature change takes place in a single phase of matter

106. SYMBOLS IN REACTIONS (s) - solid (g) - gas (l) - liquid (aq) - aqueous�.. dissolved in water See board for other symbols

107. 3 Phases of Matter The balance between attractive forces and kinetic energy determines the phase of matter High kinetic energy (KE) and low attractive forces= gas Low KE and high attractive forces=solid IntermediateKE and attractive forces=liquid

108. SOLIDS Two Main Types of Solids: 1. Crystalline 2. Amorphous CRYSTAL: the atoms, ions,or molecules are arranged in an orderly, repeating, 3-dimensional pattern (crystal lattice)

109. Allotrope Allotropes are substances with the same elemental composition, but different geometric arrangement. Carbon has 4 allotropes: a. diamond-formed under tremendous pressure b.graphite- more loosely packed c. soot- randomly bonded (amorphous form)d. buckey ball

110. AMORPHOUS SOLID An Amorphous solid lacks an ordered , internal structure. Atoms, ions, or molecules are arranged randomly. Generally, these substances are �super-cooled�-there is not enough time for the particles to arrange themselves in a pattern Examples: Rubber, glass, plastics, polymers

111. KINETIC MOLECULAR THEORY OF GASES 1. Gases are made up of very small particles called atoms or molecules 2. Gas particles are separated by large distances (low density) 3. The particles are in constant, random, straight-line motion undergoing thousands of collisions per second

112. Kinetic Molecular Theory of Gases 4. Collisions are perfectly �ELASTIC� -total kinetic energy remains constant 5. Gases exert a pressure due to the �collisions� on each other .

113. Collision Theory When gas particles collide they exert a pressure on their container Temperature is the measure of the Average Kinetic energy of the gas particles Demonstration There are 4 main properties of gases that determine their physical behavior .

114. 4 Main Properties of Gases Pressure = Force / Area Temperature : the average kinetic energy of particles Volume : space occupied by matter Amount of gas : mass (g) or (moles)

115. Effect of Pressure on Volume of a Gas

116. Increasing the Temperature of Gas at Constant Pressure. What Happens to Volume?

117. Effect of Increasing the Amount of Gas Particles

118. Same Temperature, Volume,Pressure, and Amount of GasWhat is different?

119. Lewis Structures

120. General Rules for Drawing Lewis Structures 1. Count total # of dots (valence electrons) in the structure 2. Spatially arrange the atoms. (More than two atoms - locate the central atom) 3. Try to obtain either 2 or 8 dots (valence electrons) around each atom 4. If single bonds don�t work, try double bonds, then triple bonds.

122. Figure 10.47The Phase Diagram for Water

123. 3 Factors Affecting the Rate of Solution (How fast solute will dissolve in the solvent) 1. Stirring or agitation (more solute/solvent contact at a faster rate) 2. Smaller particles (increases surface area of solute and therefore there is more solute/solvent contact at a faster rate) 3. Increase Temperature (Increases the kinetic energy and faster rate of contact between the solute/solvent particles)

124. The Solubilities of Solids as a Function of Temperature a.What is the solubility of sugar at 50 degrees Celsius? -_____ b. Which solute is least affected by an increase in temp? c. Generally, as Temp increases Solubility______

125. 260 grams NaBr Increases

126. Affect of Temperature on Solubilities of Gases in Solution Example: Opened Carbonated Drinks getting warm