Chapter 2 — A Review of the Chemistry Basics

1
Chapter 2 A Review of the Chemistry Basics

2
Atoms, Elements, Compounds and Molecules

• Atoms are the basic building blocks of matter.
• Atoms are the smallest particle of an element
  that canenter into a chemical reaction.
• Elements are comprised of a single kind of atom.
• An element is defined by its number of
  protons (Z).
• Protons have positive charges and a mass of 1
  atomic mass unit (amu). One amu 1.66 x 10-27
  kg.
• Neutrons have no charge (electrically neutral)
  and a mass of 1 atomic mass unit (amu).
• Electrons have negative charges and very small
  (negligible) mass.
• Some material taken from Chemistry, 7th Edition,
  Zumdahl Zumdahl.

3
To Speak the Language

• Atoms, Elements, Compounds and Molecules
• Atomic Names
• Inorganic Nomenclature
• Ionic Compounds
• Molecular Compounds
• Organic Nomenclature

4
Atoms, Elements, Compounds and Molecules

• Compounds are comprised of more than one kind of
  atom in a fixed ratio by mass.
• Molecules (aka molecular compounds) are groups of
  atoms chemically bonded together (covalent bonds)
  into a discrete unit. Molecules are electrically
  neutral (no net charge).
• A substance is matter that has a definite
  composition and constant properties.
• Some material taken from Introductory Chemistry
  Concepts and Connections, 5th Edition, C.H.
  Corwin.

5
Atoms, Elements, Compounds and Molecules

• Elements cannot be broken down or decomposed into
  simpler substances by chemical or physical means,
  whereas compounds can be broken down into
  elements by chemical processes.
• Can elements be molecules?
• Yes, Elements made up of 2 or more equal atoms
  form a molecule.
• Example Oxygen (O2) consists of 2 oxygen atoms
  bonded together to form an oxygen molecule.
  Oxygen is an element and a molecule.

6
Atoms, Elements, Compounds, and Molecules

• Chemical properties are characteristics that
  describe the chemical reactivity of a substance.
  Chemical changes result in the formation of
  chemically different substances (e.g., run an
  electric current thru water which contains NaSO4,
  the water will separate into hydrogen gas and
  oxygen gas). This is a chemical change from one
  substance (liquid water) into 2 new chemical
  substances (hydrogen gas and oxygen gas).
• Physical properties do not describe the chemical
  reactivity of a substance. A substance can
  display physical properties without a change in
  composition. Physical changes occur without
  changing the chemical make up of the substance
  (e.g., ice cube melts into water is a physical
  change from water in the solid form to water in
  the liquid form but IT IS STILL WATER).
• Some material taken from General Chemistry
  Principles and Modern Applications, 9th Edition,
  Petrucci, Harwood, Herring, and Madura.

7
Physical Properties

• A physical property can be observed or measured
  without changing the chemical makeup of the
  substance.
• A physical change occurs without producing a new
  chemical substance

8
Intensive and Extensive Properties

• Physical Properties have 2 categories
• Intensive properties are integral to the material
  and do not depend on the quantity of material
  (e.g., color).
• Extensive properties depend on the amount of
  material (volume, mass).

9
Atomic Number and Atomic Weight

• The atomic number (Z) of an atom is the number of
  protons in the nucleus (identity of the atom,
  e.g., carbon has z6, therefore, all carbons have
  6 protons, and all atoms having 6 protons are
  carbon atoms).
• The neutron number (N) is the number of neutrons
  in the nucleus.
• The mass number (A) of an atom is the sum of the
  proton number (or atomic number) and the neutron
  number (N). (AZN).
• The atomic mass (atomic mass unit or amu) or
  atomic weight is the average mass of an atom in a
  natural sample of the element.

10
Atomic Symbols

• The mass number is frequently written as a
  superscript and the atomic number as a subscript
  (the highest number is the atomic mass and the
  lowest number is the atomic number)

mass (A) (protonsneutrons)
atomic (Z) (protons)
Example Technetium-99 is used as a radioactive
tracer in nuclear medicine 99 43 Tc has 43
protons and 56 neutrons since 99-4356
11
Mass Spectroscopy
(series of charged metal plates, pulls and
pushes)
(gas sample bombarded with high energy
electrons, sample becomes charged since 1
electron is knocked off)
(external magnets push the molecular ion into
curved path)
(high temp. and low press. make sample a gas)
(the amount of deflection or curvature depends
on the mass of the molecular ion if too light or
too heavy it crashes into walls)
(only ions of the correct mass curve around the
bend and reach a detector)
(by manipulating the strength of external magnets
you can scan to find the atomic mass of sample)
12
Isotopes

• When J. J. Thompson first noticed a small
  signal at 22 amu in the mass spectrum (MS) of
  neon, he first assumed it was an impuritybut it
  turned out to be an isotope. He discovered the
  electron, the isotopes, and invented mass
  spectrometer.
• Isotopes have the same atomic number but
  different atomic weights (same Z different A)
• Isotopes have same number of protons, different
  number of neutrons
• How does this compare with Daltons theory?

13
Daltons Theory Background

• Law of Conservation of Mass No detectable
  change in the total mass occurs during a chemical
  reaction. During a chemical change the components
  of a system are neither created nor destroyed but
  they recombine into new substances. Antoine
  Lavoisier.
• Law of Definite Proportions Different samples
  of a pure compound always contain the same
  elements in the same proportion by mass. Example
  A sample of water taken from any source always
  contains 11.2 hydrogen and 88.8 oxygen by mass.
  Joseph Proust.
• Some material taken from General Chemistry
  Principles and Modern Applications, 9th Edition,
  Petrucci, Harwood, Herring, and Madura.

9th
14
Daltons Theory

• Dalton proposed the following three hypotheses to
  explain the Laws of Conservation of Mass and
  Definite Proportions

15
Daltons Theory

• Each element is composed of tiny, indivisible
  particles called atoms, which are identical for
  that element but are different (particularly
  their masses and chemical properties) from atoms
  of other elements

16
Daltons Theory

• Chemical combination is simply the bonding of a
  definite, small whole number of atoms of each of
  the combining elements to make one molecule of
  the formed compound. A given compound always has
  the same relative numbers and types of atoms.
  Compounds are formed by bonding atoms together in
  a fixed ratio.

17
Daltons Theory

• No atoms are gained, lost, or changed in identity
  during a chemical reaction they are just
  rearranged (recombined) to produce new substances.

18
Average Atomic Weights

• The atomic weights listed in the periodic table
  are weighted averages of the atomic masses of the
  naturally occurring isotopes of that element

19
Average Atomic Weights

• For example (chlorine)
• 75.77 of Cl has a mass of 34.96885 amu
• 24.23 of Cl has a mass of 36.96590 amu
• 0.7577 x 34.96885 26.49589
• 0.2423 x 36.96590 8.956837
• 35.4527 35.45 amu (listed in
  periodic table)

weighted averages of the atomic masses of the
naturally occurring isotopes of chlorine
20
Dmitri Mendeleev 18341907

• Russian chemist who constructed a periodic
  table of the elements, emphasizing that chemical
  and physical properties are repeated in a
  predictable way
• http//www.chemistry.co.nz/mendeleev.htm

21
Mendeleevs Periodic Table
Scandium
Ga
Ge
http//pearl1.lanl.gov/periodic/mendeleev.htm
22
Mendeleevs Periodic Table

• Gallium, scandium, and germanium were unknown
  when Mendeleev published his table in 1872.
  Mendeleev correctly predicted the existence and
  properties of these elements from gaps in his
  periodic table.
• Mendeleevs table lists the elements in order of
  atomic mass (A) instead of atomic number (Z)

23
The Modern Periodic Table

• The rows are called periods and increase by
  atomic number (increase by 1 proton and 1
  electron at a time).
• Periods represent adding electrons to quantum
  energy levels in the atom (electron shells, pages
  38-39 Textbook).
• Atoms (elements) at the end of a period each have
  an electron shell filled to its capacity with
  electrons (Noble gases only have minimal tendency
  to react very stable).
• The columns are called groups or families
• Elements within a family (group) have similar
  chemical and physical properties.
• Periodic Law The properties of elements are
  periodic functions of their atomic numbers.

24
Periodic Table
25
Information from the Table
1 H 1.00797
17 Cl 35.4527
9 F 18.9984
26
Metals, Nonmetals and Metalloids
27
Metals Metalloids NonMetals
28
Solids, Liquids and Gases
29
Representative (high-rise), Transition
(connection) and Inner Transition (footnote)
Elements
30
Overview of Matter (Pure Substances and Mixtures)
Pure Substances (cannot be physically separated)
Mixtures (can be physically separated)
31
Some Common Elements (Intensive Properties)
32
Some Common Elements (Intensive Properties)
33
Some Common Elements
34
Molecular Substances (only non-metals)

• A molecule is a group of atoms chemically bonded
  together into a discrete unit
• These types of bonds are called covalent bonds
  (sharing of electrons).
• The molecular formula of a substance gives the
  number of each kind of atom in the molecule
• Compounds composed of nonmetals tend to be
  molecular
• A helpful note Covalent Close on the periodic
  table)

35
Molecular Elements

• Diatomic
• H2, N2, O2, F2, Cl2, Br2, I2
• Tetratomic
• P4
• Octatomic
• S8

36
Ionic Compounds (usually a metal and a non-metal,
usually called salts)

• Ionic compounds are held together by ionic bonds,
  or the attraction of oppositely charged ions
• In the solid state, ionic compounds form
  crystalline lattices Most solids have periodic
  arrays of atoms which form a crystal lattice.
  Amorphous solids and glasses are exceptions.
• Cations (positive ions) are attracted to all the
  neighboring anions (negative ions), not just one
  (and viceversa).
• There are no discrete ionic molecules since one
  cannot identify any anion that is associated with
  a particular cation.

37
Ions

• An ion is an atom or group of atoms with a charge
• Ionic compounds are ions held together by ionic
  bonds
• Cations are positively charged ions
• Cations are formed by loss of electrons
• Metals tend to form cations
• Anions are negatively charged ions
• Anions are formed by gain of electrons
• Nonmetals tend to form anions

38
Monatomic Cations of Representative Metals

• For representative cations, the ionic charge
  equals the group number (vertical family) because
  the group number is equal to the number of
  electrons in the highest energy, partially filled
  electron shell (valence shell).
• Sodium metal has 11 electrons (see periodic
  table). One of these electrons will be given up
  to achieve noble gas look (the neon look)
• Na
• Mg2
• Al3
• To name these cations, name the element and add
  cation
• Sodium cation
• Magnesium cation
• Aluminum cation

39
Monatomic Anions of Representative NonMetals

• For representative anions, the ionic charge
  equals the group number minus 8
• To name monatomic anions, add the suffix ide to
  the stem name
• Cl chloride ion (group 7A
  therefore, 7-8-1) ready to accept 1 electron
• S2 sulfide ion (group 6A
  therefore, 6-8-2) ready to accept 2 electrons
• N3 nitride ion (group 5A
  therefore, 5-8-3) ready to accept 3 electrons
• O2 oxide ion (group 6A
  therefore, 6-8-2) ready to accept 2 electrons

40
Transition Metal Cations

• Most transition metals form more than 1 cation

Older system lower charged ion as the ous ion
(e.g., cuprous) and the higher charged ion as the
ic (e.g., cupric) ion
41
Naming Transition Metal Cations

• If the transition metal forms only one cation,
  name them like representative metal cations
• If the transition metal forms more than one
  cation, name the metal and give the charge in
  Roman numerals in parentheses
• e.g., Fe3 is the iron (III) ion

42
Naming Transition Metal Cations

• The older system for naming transition metals is
  to name the lower charged ion as the ous ion
  and the higher charged ion as the icion
• Cr2 is the chromous ion
• Cr3 is the chromic ion
• Use the Latin stem name if the ion name becomes
  clumsy
• Fe2 is the ferrous ion
• Fe3 is the ferric ion

43
Polyatomic Ions (from 2 or more nonmetal atoms
bonded together (ionic bond) with a resulting net
electrical charge)
44
Formulas of Ionic Compounds

• The net charge on a formula unit must be zero
• Therefore
• S () charges S () charges
• Na and
  Cl ? NaCl (sodium chloride)
• Al3 and
  O2 ? Al2O3 (aluminum oxide)
• Ca2 and
  O2 ?CaO (calcium oxide)

45
Naming Ionic Compounds

• Name the cation then the anion
• K2O
• (potassium oxide)
• Li2CO3
• (lithium carbonate)
• K2SO4
• (potassium sulfate)

46
Naming Ionic Compounds

• For transition metal cations, indicate the charge
  on the cation
• 2 -2
• FeSO4
• 3 -2
• Fe2(SO4)3

47
Formulas from Names

• What are the formulas of these compounds?
• calcium sulfide
• 2 -2
• CaS
• chromium (III) acetate
• 3 -1
• Cr(C2H3O2)3
• plumbous nitrate
• 2 -1
• Pb(NO3)2

48
Naming Molecular Compounds

• Molecular compounds made up of discrete units
  (molecules) and usually consist of a small number
  of nonmetal atoms held together by covalent
  bonds.
• Name each element
• Indicate how many of each element is present with
  a prefix multiplier
• mono 1 di 2 tri 3 tetra 4 penta 5
  hexa 6
• Add the suffix ide to the last element

Some material taken from General Chemistry
Principles and Modern Applications, 9th Edition,
Petrucci, Harwood, Herring, and Madura.
49
Naming Molecular Compounds Examples

• HCl
• hydrogen chloride
• NI3
• nitrogen triiodide
• N2O4
• dinitrogen tetraoxide

50
Molecular Compounds Common Names

• These compounds have common (nonsystematic
  names)
• Water (H2O)
• Ammonia (NH3)
• Methane (CH4)
• Ethane (C2H6)
• Propane (C3H8)
• Nitrous oxide (N2O)

51
Hydrates

• Some ionic compounds incorporate a fixed number
  of water molecules into their formula unit
• Naming hydrates only makes sense when you are
  dealing with solid reagents
• The number of waters is indicated with a
  multiplier number
• CuSO4 5 H2O copper (II) sulphate pentahydrate
• CaSO4 ½ H2O calcium sulfate sesquihydrate

52
Desiccants

• A desiccant is the anhydrous form of a compound
  that has a strong tendency to form a hydrate, and
  is used to scavenge the last traces of water from
  a system.
• NB not all desiccants form stoichiometric
  hydrates
• The most commonly used desiccant is silica gel
  (SiO2)
• Addition of water to a desiccant is a reversible
  process, so saturated desiccants can be used as
  moisturizers
• NB nota bene Latin for note well