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Chemistry 112. Overview of Chapters 1-4. 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.

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chemistry 112

Chemistry 112

Overview of Chapters 1-4

chapter 1 highlights
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.
chapter 1 highlights cont
Chapter 1 Highlights (cont)
  • 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 cont1
Chapter 1 Highlights (cont)
  • The arrangement of electrons in the orbitals is called the electronic configuration and determines the chemistry of an atom.
slide7
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.
slide8
The Building Blocks of Matter
    • Atoms
      • Smallest representative units of the elements.
    • Compounds
      • Different atoms linked together; e.g., H2O.
slide9
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.
slide10
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.
slide12
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
slide13
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.
slide15
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.
slide17
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.
slide20
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.
chapter 2 highlights
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.
chapter 2 highlights cont
Chapter 2 Highlights (cont)
  • 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 cont1
Chapter 2 Highlights (cont)
  • 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.
slide25
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.
slide29
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-.
slide31
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.
slide32
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.
slide35
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.
chapter 3 highlights
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.
chapter 3 highlights1
Chapter 3 Highlights
  • 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).
slide39
Na + Cl NaCl
  • 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.

Reactants Product

slide40
Incorrect

H2 + O2 H2O

2 H2 + O2 2 H2O

Correct

  • 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.
slide41
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.
slide42
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 1023 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.
slide44
Stoichiometry Calculations
    • The units of molar mass are grams/mole.
    • Moles x molar mass = mass.
      • Example: 2.0 mol CO2 x 44 g/mol = 88 g CO2
    • Mass/molar mass= moles.
      • Example: 132 g CO2 / 44 g/mol = 3.0 mol CO2
slide45
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.
slide46
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.
slide48
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.
chapter 4 highlights
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

chapter 4 highlights cont
Chapter 4 Highlights (cont)
  • 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 acidic; pH > 7.0 is basic; pH = 7.0 is neutral.
  • Acids react with bases in neutralization reactions.
slide52
States of Matter
    • Review of Types of Bonds
      • Chemical bonds (intramolecular forces) hold atoms together.
      • The three types of chemical bonds are ionic, polar covalent, and nonpolar covalent.
      • Intermolecular forces hold molecules together.
chapter outline
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.
slide56
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.
slide58
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.
slide60
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 H2O.
      • Hydrogen bonds are typically quite strong.
slide63
Forming Solutions
    • Like dissolves like
      • Ionic solutes often dissolve in polar solvents;e.g., NaCl dissolves in H2O.
      • Polar solutes generally dissolve in polar solvents; e.g., NH3 in H2O.
      • Nonpolar solutes generally do not dissolve well in polar solvents; e.g., oil in H2O.
slide65
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”).
slide67
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).
slide68
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.
slide69
Acid-Base Chemistry (cont)
    • The pH Scale: a measure of acidity
slide70
Acid-Base Chemistry (cont)
    • Acid-Base Indicators
      • Molecular sensors of H+.

H+

slide71
Acid-Base Chemistry (cont)
    • Neutralization Reactions: equal molar amounts of an acid and a base react to form a neutral solution.

HCl + NaOH NaCl + H2O

slide72
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 HAto make more A- and H2O
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