Chemistry 112
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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

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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.

  • The different types of atoms are called elements, which are arranged systematically in the periodic table.

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 (cont)

  • The arrangement of electrons in the orbitals is called the electronic configuration and determines the chemistry of an atom.

Types of Matter

States of Matter

  • 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.

  • The Building Blocks of Matter

    • Atoms

      • Smallest representative units of the elements.

    • Compounds

      • Different atoms linked together; e.g., H2O.

  • 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 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

  • 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 (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.


  • 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.

The Rutherford Experiment

  • 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.

The Solar System Model

Electromagnetic Radiation

  • 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.

The Orbital Model:Electronic Configurations

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)

  • 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)

  • 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.

  • 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.

ElectronicConfiguration of Noble Gases

Types of Compounds

Lewis Dot Structures

  • 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 Compounds

  • 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.

  • 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.


Polar vs. Nonpolar Bonds

  • 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.

Drawing Molecules

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 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).

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


H2 + O2 H2O

2 H2 + O2 2 H2O


  • 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)

    • 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.

  • 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.

The Mole

  • 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

  • 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.

  • 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 (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

  • 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)

  • 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.

  • 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.

Review of Types of Bonds

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.

Changes of State

  • 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.

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.

Dipole-Dipole Interactions

  • 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.

Hydrogen Bonds in Water


  • 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.

NaCl Dissolving in H2O

  • 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”).

Emulsification with Soap

  • 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).

  • 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 (cont)

    • The pH Scale: a measure of acidity

  • Acid-Base Chemistry (cont)

    • Acid-Base Indicators

      • Molecular sensors of H+.


  • 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

  • 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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