1. Chemistry 112 Overview of Chapters 1-4

2. Chapter 1 Highlights  Chemistry is the study of matter, the physical substance of all materials.  The building blocks of matter are atoms, which combine to form compounds.  The different types of atoms are called elements, which are arranged systematically in the periodic table.  Chemistry is the study of matter, the physical substance of all materials.  The building blocks of matter are atoms, which combine to form compounds.  The different types of atoms are called elements, which are arranged systematically in the periodic table.

3.  Atoms are composed of protons, neutrons, and electrons.  All atoms of the same element contain the same number of protons (and electrons) but may vary in the number of neutrons.  The protons and neutrons are found inside the tiny but dense nucleus, whereas the electrons are found in orbitals outside the nucleus.  Atoms are composed of protons, neutrons, and electrons.  All atoms of the same element contain the same number of protons (and electrons) but may vary in the number of neutrons.  The protons and neutrons are found inside the tiny but dense nucleus, whereas the electrons are found in orbitals outside the nucleus. Chapter 1 Highlights (cont)

4.  The arrangement of electrons in the orbitals is called the electronic configuration and determines the chemistry of an atom. Chapter 1 Highlights (cont)

5. Types of Matter

6. States of Matter

7.  Chemistry and Matter  Physical Changes versus Chemical Changes  Physical changes involve changes in appearance (i.e., changes in state such as melting).  Chemical changes result in new substances.  Chemistry and Matter  Physical Changes versus Chemical Changes  Physical changes involve changes in appearance (i.e., changes in state such as melting).  Chemical changes result in new substances.

8.  The Building Blocks of Matter  Atoms  Smallest representative units of the elements.  Compounds  Different atoms linked together; e.g., H 2 O.  The Building Blocks of Matter  Atoms  Smallest representative units of the elements.  Compounds  Different atoms linked together; e.g., H 2 O.

9.  The Building Blocks of Matter (cont)  Dalton’s Atomic Theory  All matter is composed of indivisible atoms.  All atoms of one element are identical to each other but different than the atoms of other elements.  Compounds are formed when atoms of different elements combine in whole number ratios.  Atoms are rearranged during chemical reactions but atoms cannot be created or destroyed.  The Building Blocks of Matter (cont)  Dalton’s Atomic Theory  All matter is composed of indivisible atoms.  All atoms of one element are identical to each other but different than the atoms of other elements.  Compounds are formed when atoms of different elements combine in whole number ratios.  Atoms are rearranged during chemical reactions but atoms cannot be created or destroyed.

10.  The Periodic Table  Used to organize the elements by recurring chemical properties.  Elements in the same vertical column of the periodic table have similar chemical properties and are said to be in the same group or family.  The Periodic Table  Used to organize the elements by recurring chemical properties.  Elements in the same vertical column of the periodic table have similar chemical properties and are said to be in the same group or family.

11. The Periodic Table

12.  The Atom  Components  Positive protons, negative electrons, and neutral neutrons  Atomic Number  The number of protons in an atom, which determines what element it is  Mass Number  Number of protons + the number of neutrons  The Atom  Components  Positive protons, negative electrons, and neutral neutrons  Atomic Number  The number of protons in an atom, which determines what element it is  Mass Number  Number of protons + the number of neutrons

13.  The Atom (cont)  Isotopes  Isotopes of the same element have the same number of protons but differ in the number of neutrons.  Atomic Mass  The atomic mass for each element on the periodic table reflects the relative abundance of each isotope in nature.  The Atom (cont)  Isotopes  Isotopes of the same element have the same number of protons but differ in the number of neutrons.  Atomic Mass  The atomic mass for each element on the periodic table reflects the relative abundance of each isotope in nature.

14. Isotopes

15.  Models of the Atom  The Plum Pudding Model  Electrons are embedded in a sphere of positive charge.  The Nuclear Model  All of the positive charge is in a tiny central nucleus with electrons outside the nucleus.  This model was developed by Rutherford after his landmark experiment.  Models of the Atom  The Plum Pudding Model  Electrons are embedded in a sphere of positive charge.  The Nuclear Model  All of the positive charge is in a tiny central nucleus with electrons outside the nucleus.  This model was developed by Rutherford after his landmark experiment.

16. The Rutherford Experiment

17.  Models of the Atom (continued)  Bohr’s Solar System Model  Electrons circle the nucleus in orbits, which are also called energy levels.  An electron can “jump” from a lower energy level to a higher one upon absorbing energy, creating an excited state.  The concept of energy levels accounts for the emission of distinct wavelengths of electromagnetic radiation during flame tests.  Models of the Atom (continued)  Bohr’s Solar System Model  Electrons circle the nucleus in orbits, which are also called energy levels.  An electron can “jump” from a lower energy level to a higher one upon absorbing energy, creating an excited state.  The concept of energy levels accounts for the emission of distinct wavelengths of electromagnetic radiation during flame tests.

18. The Solar System Model

19. Electromagnetic Radiation

20.  Models of the Atom (continued)  The Modern Model  Orbits are replaced with orbitals, volumes of space where the electrons can be found.  The arrangement of electrons in the orbitals is the electronic configuration of an atom, which determines the chemistry of an atom.  Models of the Atom (continued)  The Modern Model  Orbits are replaced with orbitals, volumes of space where the electrons can be found.  The arrangement of electrons in the orbitals is the electronic configuration of an atom, which determines the chemistry of an atom.

21. The Orbital Model: Electronic Configurations

22. Chapter 2 Highlights  Having eight valence electrons is particularly desirable (“the octet rule”).  Atoms form bonds with other atoms to satisfy the octet rule.  The two major types of chemical bonds are ionic and covalent.  Having eight valence electrons is particularly desirable (“the octet rule”).  Atoms form bonds with other atoms to satisfy the octet rule.  The two major types of chemical bonds are ionic and covalent.

23.  Electronegativity is the ability to attract shared electrons.  The type of bond formed between two atoms depends on their difference in electronegativity.  Ionic bonds form between atoms with a large difference in electronegativity (generally a metal and a nonmetal).  Electronegativity is the ability to attract shared electrons.  The type of bond formed between two atoms depends on their difference in electronegativity.  Ionic bonds form between atoms with a large difference in electronegativity (generally a metal and a nonmetal). Chapter 2 Highlights (cont)

24.  Nonpolar covalent bonds form between atoms with little difference in electronegativity (generally two nonmetals).  Polar covalent bonds form between atoms with intermediate difference in electronegativity.  There are many ways to depict molecules.  Nonpolar covalent bonds form between atoms with little difference in electronegativity (generally two nonmetals).  Polar covalent bonds form between atoms with intermediate difference in electronegativity.  There are many ways to depict molecules. Chapter 2 Highlights (cont)

25.  The Octet Rule  Atoms with eight valence electrons are particularly stable, an observation called the octet rule.  Atoms form bonds with other atoms to achieve a valence octet.  The Octet Rule  Atoms with eight valence electrons are particularly stable, an observation called the octet rule.  Atoms form bonds with other atoms to achieve a valence octet.

26. Electronic Configuration of Noble Gases

27. Types of Compounds

28. Lewis Dot Structures

29.  Ionic Bonds  Ionic compounds result from the loss of electrons by one atom (usually a metal) and the gain of electrons by another atom (usually a nonmetal).  Ionic bonds arise from the attraction between particles with opposite charges (electrostatic forces); e.g., Na + Cl -.  Ionic Bonds  Ionic compounds result from the loss of electrons by one atom (usually a metal) and the gain of electrons by another atom (usually a nonmetal).  Ionic bonds arise from the attraction between particles with opposite charges (electrostatic forces); e.g., Na + Cl -.

30. Ionic Compounds

31.  Covalent Bonds  Covalent bonds are formed when two atoms share one or more electron pairs.  When two atoms share one pair of electrons, the result is a single bond.  Two shared pairs of electrons is a double bond; three is a triple bond.  Covalent Bonds  Covalent bonds are formed when two atoms share one or more electron pairs.  When two atoms share one pair of electrons, the result is a single bond.  Two shared pairs of electrons is a double bond; three is a triple bond.

32.  Equal Sharing versus Unequal Sharing  When two different kinds of atoms are bonded, the electrons are usually shared unequally.  When a bond exists between two identical kinds of atoms, the electrons are shared equally.  An atom with greater electronegativity has a greater ability to attract shared electrons.  Equal Sharing versus Unequal Sharing  When two different kinds of atoms are bonded, the electrons are usually shared unequally.  When a bond exists between two identical kinds of atoms, the electrons are shared equally.  An atom with greater electronegativity has a greater ability to attract shared electrons.

33. Electronegativity

34. Polar vs. Nonpolar Bonds

35.  Representing Structures  In a structural formula, atoms are represented by chemical symbols, and bonds are represented by lines.  In a line drawing, any point where lines connect or terminate is understood to be a carbon atom with sufficient bonded hydrogen atoms to achieve the four bonds necessary for carbon.  Representing Structures  In a structural formula, atoms are represented by chemical symbols, and bonds are represented by lines.  In a line drawing, any point where lines connect or terminate is understood to be a carbon atom with sufficient bonded hydrogen atoms to achieve the four bonds necessary for carbon.

36. Drawing Molecules

37. Chapter 3 Highlights  Reaction equations have with the initial materials (reactants) on the left, followed by a reaction arrow pointing from left to right, and the final materials (products) on the right.  A balanced equation has the same number and kinds of atoms on both sides of the equation.  Reaction equations have with the initial materials (reactants) on the left, followed by a reaction arrow pointing from left to right, and the final materials (products) on the right.  A balanced equation has the same number and kinds of atoms on both sides of the equation.

38.  The relationship between the amounts of reactants and products is the stoichiometry, which comes from a balanced reaction equation.  The SI unit for measuring atoms and molecules is the mole.  In an oxidation-reduction reaction, electrons are transferred from one material (the substance that is oxidized) to another material (the substance that is reduced).  The relationship between the amounts of reactants and products is the stoichiometry, which comes from a balanced reaction equation.  The SI unit for measuring atoms and molecules is the mole.  In an oxidation-reduction reaction, electrons are transferred from one material (the substance that is oxidized) to another material (the substance that is reduced). Chapter 3 Highlights

39.  Balanced Reaction Equations  Writing a Chemical Reaction  The starting materials, the reactants, are written on the left.  The materials that are produced, the products, are written on the right.  Reactants are separated from products by a horizontal arrow pointing from left to right.  Balanced Reaction Equations  Writing a Chemical Reaction  The starting materials, the reactants, are written on the left.  The materials that are produced, the products, are written on the right.  Reactants are separated from products by a horizontal arrow pointing from left to right. Na + Cl NaCl Reactants Product

40.  Balanced Reaction Equations (cont)  Balancing the Equation  The law of conservation of matter states that matter can neither be created nor destroyed in a chemical reaction.  The number and kind of atoms on the left- hand side of an equation must be equal to the number and kind of atoms on the right.  Balanced Reaction Equations (cont)  Balancing the Equation  The law of conservation of matter states that matter can neither be created nor destroyed in a chemical reaction.  The number and kind of atoms on the left- hand side of an equation must be equal to the number and kind of atoms on the right. H 2 + O 2 H 2 OIncorrect 2 H 2 + O 2 2 H 2 OCorrect

41.  Balanced Reaction Equations (cont)  Stoichiometry  The stoichiometry of a chemical reaction is the relationship between the number of molecules of the reactants and products in the balanced reaction equation.  A reactant present in insufficient amounts is the limiting reagent.  Balanced Reaction Equations (cont)  Stoichiometry  The stoichiometry of a chemical reaction is the relationship between the number of molecules of the reactants and products in the balanced reaction equation.  A reactant present in insufficient amounts is the limiting reagent.

42.  The Mole  The mole is the SI unit of measure to describe the amount of matter that is present.  One mole is equal to 6.02 x 10 23 particles (Avogadro’s number).  One mole of an element has a mass that is equal to the atomic mass of that element in grams.  One mole of a compound has a mass that is equal to the molecular/formula mass of that compound in grams.  The Mole  The mole is the SI unit of measure to describe the amount of matter that is present.  One mole is equal to 6.02 x 10 23 particles (Avogadro’s number).  One mole of an element has a mass that is equal to the atomic mass of that element in grams.  One mole of a compound has a mass that is equal to the molecular/formula mass of that compound in grams.

43. The Mole

44.  Stoichiometry Calculations  The units of molar mass are grams/mole.  Moles x molar mass = mass.  Example: 2.0 mol CO 2 x 44 g/mol = 88 g CO 2  Mass/molar mass= moles.  Example: 132 g CO 2 / 44 g/mol = 3.0 mol CO 2  Stoichiometry Calculations  The units of molar mass are grams/mole.  Moles x molar mass = mass.  Example: 2.0 mol CO 2 x 44 g/mol = 88 g CO 2  Mass/molar mass= moles.  Example: 132 g CO 2 / 44 g/mol = 3.0 mol CO 2

45.  Stoichiometry Calculations  The expected mass of a product or reactant can be calculated for any reaction by using the balanced equation and the molar mass.  Stoichiometry Calculations  The expected mass of a product or reactant can be calculated for any reaction by using the balanced equation and the molar mass.

46.  Oxidation-Reduction Reactions  Defined  Oxidation-reduction (“redox”) reactions involve the transfer of electrons from one substance to another.  Oxidized substances lose electrons and reduced substances gain electrons.  Oxidation-Reduction Reactions  Defined  Oxidation-reduction (“redox”) reactions involve the transfer of electrons from one substance to another.  Oxidized substances lose electrons and reduced substances gain electrons.

47. Oxidation-Reduction

48.  Oxidation-Reduction Reactions (cont)  The Chemistry of Batteries  Combining a readily oxidized substance with an easily reduced substance can create a battery.  The oxidized material is the anode and the reduced material is the cathode of the battery.  Oxidation-Reduction Reactions (cont)  The Chemistry of Batteries  Combining a readily oxidized substance with an easily reduced substance can create a battery.  The oxidized material is the anode and the reduced material is the cathode of the battery.

49. Batteries

50. Chapter 4 Highlights  Intermolecular forces hold the molecules of a material together.  Stronger intermolecular forces lead to higher melting and boiling temperatures.  The relative strengths of intermolecular forces generally follow the trend: hydrogen bonds > dipole-dipole interactions > London forces  Intermolecular forces hold the molecules of a material together.  Stronger intermolecular forces lead to higher melting and boiling temperatures.  The relative strengths of intermolecular forces generally follow the trend: hydrogen bonds > dipole-dipole interactions > London forces

51.  Like dissolves like. That is, polar solutes dissolve in polar solvents.  Acids are proton (H + ) donors; bases are proton acceptors that produce OH - in solution.  The pH measures the acidity of a solution: pH 7.0 is basic; pH = 7.0 is neutral.  Acids react with bases in neutralization reactions.  Like dissolves like. That is, polar solutes dissolve in polar solvents.  Acids are proton (H + ) donors; bases are proton acceptors that produce OH - in solution.  The pH measures the acidity of a solution: pH 7.0 is basic; pH = 7.0 is neutral.  Acids react with bases in neutralization reactions. Chapter 4 Highlights (cont)

52.  States of Matter  Review of Types of Bonds 1.Chemical bonds (intramolecular forces) hold atoms together. 2.The three types of chemical bonds are ionic, polar covalent, and nonpolar covalent. 3.Intermolecular forces hold molecules together.  States of Matter  Review of Types of Bonds 1.Chemical bonds (intramolecular forces) hold atoms together. 2.The three types of chemical bonds are ionic, polar covalent, and nonpolar covalent. 3.Intermolecular forces hold molecules together.

53. Review of Types of Bonds

54. Chapter Outline  States of Matter (cont)  Particle Cohesion Determines Physical State  In general, the relative strengths of intermolecular forces follows the trend: gases < liquids < solids  Changes of State  Adding energy breaks intermolecular forces and causes molecules to change their state.  The stronger the intermolecular forces of a compound, the higher are the melting and boiling points.  States of Matter (cont)  Particle Cohesion Determines Physical State  In general, the relative strengths of intermolecular forces follows the trend: gases < liquids < solids  Changes of State  Adding energy breaks intermolecular forces and causes molecules to change their state.  The stronger the intermolecular forces of a compound, the higher are the melting and boiling points.

55. Changes of State

56.  Types of Intermolecular Forces within Pure Substances  London dispersion forces  A temporary dipole in one molecule can induce a dipole in a neighboring molecule.  The negative end of one temporary dipole can attract the positive end of an induced dipole; these attractions are called London dispersion forces.  London forces tend to be fairly weak.  Types of Intermolecular Forces within Pure Substances  London dispersion forces  A temporary dipole in one molecule can induce a dipole in a neighboring molecule.  The negative end of one temporary dipole can attract the positive end of an induced dipole; these attractions are called London dispersion forces.  London forces tend to be fairly weak.

57. London Dispersion Forces

58.  Types of Intermolecular Forces within Pure Substances (cont)  Dipole-dipole interactions  Dipole-dipole interactions exist between molecules with polar covalent bonds.  Dipole-dipole interactions are typically stronger than London dispersion forces.  Types of Intermolecular Forces within Pure Substances (cont)  Dipole-dipole interactions  Dipole-dipole interactions exist between molecules with polar covalent bonds.  Dipole-dipole interactions are typically stronger than London dispersion forces.

59. Dipole-Dipole Interactions

60.  Types of Intermolecular Forces within Pure Substances (cont)  Hydrogen Bonds  Hydrogen bonds are a special type of dipole-dipole interaction.  Hydrogen bonds can occur when H is bonded to one of the highly electronegative atoms N, O, or F. An example is H 2 O.  Hydrogen bonds are typically quite strong.  Types of Intermolecular Forces within Pure Substances (cont)  Hydrogen Bonds  Hydrogen bonds are a special type of dipole-dipole interaction.  Hydrogen bonds can occur when H is bonded to one of the highly electronegative atoms N, O, or F. An example is H 2 O.  Hydrogen bonds are typically quite strong.

61. Hydrogen Bonds in Water

62. Mixtures

63.  Forming Solutions  Like dissolves like  Ionic solutes often dissolve in polar solvents; e.g., NaCl dissolves in H 2 O.  Polar solutes generally dissolve in polar solvents; e.g., NH 3 in H 2 O.  Nonpolar solutes generally do not dissolve well in polar solvents; e.g., oil in H 2 O.  Forming Solutions  Like dissolves like  Ionic solutes often dissolve in polar solvents; e.g., NaCl dissolves in H 2 O.  Polar solutes generally dissolve in polar solvents; e.g., NH 3 in H 2 O.  Nonpolar solutes generally do not dissolve well in polar solvents; e.g., oil in H 2 O.

64. NaCl Dissolving in H 2 O

65.  Emulsions  Emulsifying agents are molecules that contain a polar portion and a nonpolar region.  Soap is an example of an emulsifying agent that can form a suspension of a nonpolar material in a polar solvent (an “emulsion”).  Emulsions  Emulsifying agents are molecules that contain a polar portion and a nonpolar region.  Soap is an example of an emulsifying agent that can form a suspension of a nonpolar material in a polar solvent (an “emulsion”).

66. Emulsification with Soap

67.  Measuring Amounts in Solution  Solubility  The maximum amount of a solute that dissolves in a solvent  Molarity  The amount of a solute dissolved in a solvent is its concentration.  Concentration is often measured in moles/liter, also called molarity (M).  Measuring Amounts in Solution  Solubility  The maximum amount of a solute that dissolves in a solvent  Molarity  The amount of a solute dissolved in a solvent is its concentration.  Concentration is often measured in moles/liter, also called molarity (M).

68.  Acid-Base Chemistry  Definitions of Acids and Bases  Acids turn litmus paper red; bases turn litmus paper blue.  Acids produce H + in solution; bases produce OH - in solution.  Acids are proton donors; bases are proton acceptors.  Acid-Base Chemistry  Definitions of Acids and Bases  Acids turn litmus paper red; bases turn litmus paper blue.  Acids produce H + in solution; bases produce OH - in solution.  Acids are proton donors; bases are proton acceptors.

69.  Acid-Base Chemistry (cont)  The pH Scale: a measure of acidity  Acid-Base Chemistry (cont)  The pH Scale: a measure of acidity

70.  Acid-Base Chemistry (cont)  Acid-Base Indicators  Molecular sensors of H +.  Acid-Base Chemistry (cont)  Acid-Base Indicators  Molecular sensors of H +. H+H+

71.  Acid-Base Chemistry (cont)  Neutralization Reactions: equal molar amounts of an acid and a base react to form a neutral solution.  Acid-Base Chemistry (cont)  Neutralization Reactions: equal molar amounts of an acid and a base react to form a neutral solution. HCl + NaOH NaCl + H 2 O

72.  Acid-Base Chemistry (cont)  Buffers: contain a weak acid and its conjugate base, which react with added H + or OH - to prevent pH changes.  Acid-Base Chemistry (cont)  Buffers: contain a weak acid and its conjugate base, which react with added H + or OH - to prevent pH changes. HA H + + A - Adding acid: H + reacts with A - to make more HA Adding base: OH - reacts with HA to make more A - and H 2 O