Act Chem2

1
Active Chemistry

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  should only be used for educational purposes
  (Fair Use Policy).

2
Matter

• Matter has mass and occupies space.
• It is composed of tiny particles.

3
Matter

• Matter undergoes physical and chemical changes.

• A physical change involves a change in one or
  more physical properties but no change in
  composition.
• Ex. Three physical states
• Solid
• Liquid
• Gas

4
Matter

• Matter undergoes physical and chemical changes.

• A chemical change transforms a substance into one
  or more new substances.

5
The Atom

• Atoms are tiny particles of matter.
• Atoms are composed of smaller subatomic
  particles
• The nucleus, which is at the center of the atom,
  contains protons (positively charged) and
  neutrons (uncharged).
• Electrons (negatively charged) move around the
  nucleus.

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The Atom

• Atoms of the same type form elements.

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The Periodic Table

• The periodic table lists all the elements
  discovered so far.
• Elements are arranged by categories based on
  their properties.
• The number with a letter at the top of each
  column is the Group Number
• Ex. Group 1A, 2A, 7B, 5A, etc.
• The number at the beginning of the left side of
  the rows represents the Period Number
• Ex. 1st period, 2nd period, etc.

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How to Read a Box on the Periodic Table

• Each element has a name and a symbol.
• The symbol usually consists of the first one or
  two letters of the elements name.
• Examples Oxygen O Krypton Kr
• The atomic number is the number of protons the
  atom has
• It can also represent the number of electrons and
  neutrons in a normal, neutral atom

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A. The Structure of the Atom

• Experiments by J.J. Thomson showed that atoms
  contain electrons.
• Cathode ray tube

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A. The Structure of the Atom

• Rutherfords Experiment

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A. The Structure of the Atom

• Results of the Rutherford experiment

(a) The results that the metal foil experiment
would have yielded if the plum pudding model had
been correct
(b) Actual results
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B. Introduction to the Modern Concept of Atomic
Structure

• Comparing the Parts of an Atom

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How to Read a Box on the Periodic Table

• The atomic mass represents how much matter is
  in one atom.
• Most of the mass is contained in the atoms
  nucleus

14
Metals

• Physical Properties of Metals

• Efficient conduction of heat and electricity
  (these flow easily through them)
• Malleability (can be hammered into thin sheets)
• Ductility (can be pulled into wires)
• A lustrous (shiny) appearance

15
Semimetals

• Semimetals are also known as metalloids
• Have some characteristics of both metals and
  non-metals
• Often make good semiconductors (conducts
  electricity not so readily)

16
Non-metals

• Are poor conductors of heat and electricity when
  compared to metals
• In solid form, they are dull and brittle

17
Noble Gases

• Group 8A
• Exist as gases at room temp.
• Non-metals
• 8 electrons in the outer shell Full
• Helium (He) has only 2 electrons in the outer
  shell Full
• Not reactive with other elements

18
Halogens

• Group 7A
• 7 electrons in the outer shell
• All are non-metals
• Very reactive are often bonded with elements from
  Group 1A

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Alkali Metals

• Group 1A
• Hydrogen is not a member, it is a non-metal
• 1 electron in the outer shell
• Very reactive, esp. with water
• Conduct electricity

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Alkaline Earth Metals

• Group 2A
• 2 electrons in the outer shell
• Reactive, but less than Alkali metals
• Conduct electricity

21
Transition Metals

• Groups in the middle
• Good conductors of heat and electricity.
• Some are used for jewelry.
• The transition metals are able to put up to 32
  electrons in their second to last shell.
• Can bond with many elements in a variety of
  shapes.

22
Atomic Size and Radii Trends

• Atomic radius (pl. radii) the distance between
  the outermost electrons and the nucleus of the
  atom
• The size is based on the number of electron
  shells (energy levels)
• The period number (see p. 14-15) represents the
  number of shells the atom has
• Ex. Oxygen is at period 2, so it has 2 electron
  shells

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Atomic Size and Radii Trends

• Atomic radius (pl. radii) the distance between
  the outermost electrons and the nucleus of the
  atom
• The size is based on the number of electron
  shells (energy levels)
• more shells, larger radius
• the more electrons and protons, the stronger the
  attractive force, pulling the electrons closer,
  and causing a smaller radius
• It increases moving from right to left
• It increases moving from top to bottom

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Ions

• Atoms can form ions by gaining or losing
  electrons.
• In order to reach a full outer electron shell
  (noble gas configuration), metals tend to lose
  one or more electrons to form positive ions
  called cations.
• Ex., if K loses 1 electron, it will be written as
  K which has a positive 1 charge
• Ex., if Ca loses 2 electrons, it will be written
  as Ca2 which has a positive 2 charge

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Ions

• Atoms can form ions by gaining or losing
  electrons.
• In order to reach a full outer electron shell
  (noble gas configuration), nonmetals tend to gain
  one or more electrons to form negative ions
  called anions.
• Ex., if Br gains 1 electron, it will be written
  as Br- which has a negative 1 charge
• Ex., if S gains 2 electrons, it will be written
  as S2- which has a negative 2 charge

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Ions

• The ion that a particular atom will form can be
  predicted from the periodic table.
• The atom will gain or lose electrons the easiest
  way in order to have a full outer shell.
• Ex. Beryllium in Group 2A has 2 outer electrons.
  The easiest way for it to have a full outer shell
  (2 for the first shell, 8 for the rest) is to
  lose 2 electrons. Therefore, Group 2A forms 2
  charged cations.

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Ions

• Ion Charges and the Periodic Table

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Ionic Radii Trends

• Ionic radius (pl. radii) the distance between
  the outermost electrons and the nucleus of the
  ion (see p. 37)
• Cations are always smaller than the parent atom
• Anions are always larger than the parent atom

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Ionic Radii Trends
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Ionic Radii Trends

• Ionic radius (pl. radii) the distance between
  the outermost electrons and the nucleus of the
  ion
• It increases moving from right to left
• It increases moving from top to bottom

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Electronegativity Trends

• Electronegativity (electron affinity) the
  ability of an atom in a molecule to attract
  shared electrons to itself

• In a molecule one atom attracts the electrons
  more than the other atom

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Electronegativity Trends

• Electronegativity (electron affinity) the
  relative ability of an atom in a molecule to
  attract shared electrons to itself
• It increases moving from left to right
• It increases moving from bottom to top

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Ionization Energy Trends

• Ionization energy is the energy required to
  remove the outermost (highest energy) electron
  from a neutral atom at its ground state

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Ionization Energy Trends

• Ionization energy is the energy required to
  remove the outermost (highest energy) electron
  from a neutral atom at its ground state
• It increases moving from left to right
• It increases moving from bottom to top

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Periodic Table Trends
36
Periods

• 1st Period 1 electron shell
• 2nd Period 2 electron shells
• 3rd Period 3 electron shells

• Each row is called a period
• The elements in each period have the same number
  of electron shells

37
Groups

• Group 1 1 outer electron
• Group 2 2 outer electrons
• Group 8 8 outer electrons (except for He, which
  has 2)

• Each column is called a group
• Each element in a group has the same number of
  electrons in their outer shell

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Transition Metals

• Transition Metals have slightly different rules
  for shells and valence electrons.

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Valence Electrons

• The electrons in the outer shell are called
  valence electrons which is determined by the
  group number (see p. 14-15)
• In writing Lewis dot structures we include only
  the valence electrons

C
40
Practice

• Draw valence electrons (Lewis dot structures) for
  the following elements
• sodium
• neon
• beryllium
• hydrogen
• sulfur
• potassium
• chlorine

41
Molecules

• A molecule is a particle that is made up of two
  or more atoms bonded together
• A compound is a substance made of one kind of
  molecule

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Molecules
Nitrogen gas contains N2 molecules.
Oxygen gas contains O2 molecules.
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Molecules
44
Molecules
Carbon atoms
Diamond
Graphite
Buckminsterfullerene
45
Molecules
46
Molecules

• Bonds are shared electrons (2 or more) that store
  energy
• Connect atoms together to form molecules

47
Molecular Formulas

• Steps to writing molecular formulas
• Each atom present is represented by its element
  symbol.
• The number of each type of atom is indicated by a
  subscript written to the right of the element
  symbol.
• When only one atom of a given type is present,
  the subscript 1 is not written.

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Molecular Formulas

• Steps to writing molecular formulas
• Base numbers of atoms within parentheses are
  multiplied by the subscript to the right of the
  end parentheses.
• Numbers in front of the whole formula are
  multiplied to each kind of atom.

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Molecular Formulas
Example 3Mg(NO3)2 Mg magnesium 1 x 3 3
atoms N nitrogen 1 x 2 x 3 6 atoms O
oxygen 3 x 2 x 3 18 atoms
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Chemical Reactions

• A chemical reaction is when bonds between atoms
  form or break apart
• Reactants what the reaction starts with
• Products what the reaction ends up with
• For example, 2H2 O2 gt 2H2O
• reactants products

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Indications of a Chemical Reaction
52
Chemical Reaction
53
Metallic Bonding

• Metals are made of closely packed cations (see p.
  37), surrounded by a sea of electrons (meaning
  that electrons freely move)
• This explains why metals have these properties
  (see p. 19)
• Metallic bonds are the attraction of free
  floating valence electrons to the metal cations
  (see p. 18-19, 37).

54
Covalent Bonding

• Molecules whose atoms are held together by
  covalent bonds are usually non-metals (see p.
  22-23).
• For example, CH4

55
Ionic Bonding

• Molecules whose atoms are held together by ionic
  bonds are usually a metal cation (see p. 18-19,
  37) and a non-metal anion (see p. 22-23, 37).
• For example, BeCl2
• These atoms are held by strong attractive forces
  called electrostatic forces, because they have
  opposite charges

56
Practice

• Do the following molecules have metallic, ionic,
  or covalent bonding?
• CO2
• O3
• MgBr2
• C2H2
• CrCo2
• BeF2
• CCl4
• H2CS
• KCl
• PtAu

57
Salt Crystals

• Salt crystals are ionic compounds made of cations
  and anions
• Held together by strong electrostatic forces
  between ions (see p. 55)
• Have high melting points
• As a solid, they do not conduct electricity
  because of its crystalline structure
• If melted, they conduct electricity
• For example, NaCl

58
Salt Solutions

• Salt solutions are salt crystals dissolved in
  water
• Conduct electricity

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Salt Solutions

• How to Predict the Formula of an Ionic Compound
  (see p. 57)
• Write the cation element symbol followed by the
  anion element symbol.
• The charges and numbers of the anions and cations
  in the molecule must sum to zero.
• For example magnesium and chlorine

60
Salt Solutions

• The Criss Cross Method is used to find the
  formula for ionic compounds.
• For monatomic (one atom) ions

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Salt Solutions

• For polyatomic (many atoms) ions
• Al(OH)3 Ca(NO3)2

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Practice

• Predict the formulas of the ionic compounds
  formed by the following ions
• lithium and oxygen
• calcium and sulfur
• strontium and fluorine
• magnesium and bromine

63
Practice

• Predict the formulas of the ionic compounds
  formed by the following molecules
• Barium and OH
• Calcium and NO3
• Potassium and SO42
• Oxygen and NH4

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Neutral pH

• pH means power of hydrogen and is the
  measurement of the concentration of hydroxide
  (OH) and hydrogen (H) ions in an aqueous
  solution (mixture with water).
• The amount of pH can affect chemical reactions
  (see p. 49)
• Scale is from 0 to 14

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Neutral pH

• A neutral solution has a pH of 7.
• Has an equal concentration of H and OH ions
• Ex. water and solutions
• in most living systems

pure water pH 7
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The pH Scale
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Acids

• An acid releases H ions when it dissolves in
  water.
• high H concentration
• pH less than 7
• Tastes sour
• Stings on cuts
• Releases CO2
• when reacting with
• carbonates
• Releases
• H2 when reacting
• with a metal

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Acids

• Litmus paper turns red to blue
• Strong acids dissociate (separate) fully into
  their ions and have high electric conductivity,
  weak acids do not
• Common strong acids are sulfuric acid and
  hydrochloric acid, HCl
• Common weak acids are acetic acid and citric acid

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Bases

• A base removes H ions from a solution.
• low H concentration
• pH greater than 7
• Feels slippery
• Tastes bitter
• Litmus paper
• turns blue to red
• Turns into
• water and salt
• when mixed with
• an acid

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Bases

• Strong bases dissociate (separate) fully into
  their ions and have high electric conductivity,
  weak bases do not
• Common strong bases are hydroxides like NaOH and
  Ca(OH)2
• Common weak bases are sodium bicarbonate and
  ammonia, NH3

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Exothermic Processes

• In an exothermic process, energy is released
• The evidence that you see in the following
  examples is that the surroundings get warmer
• The net heat released to the surroundings comes
  from the making and breaking of bonds during a
  chemical reaction (see p. 46-49).
• Condensation
• Freezing

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Endothermic Processes

• In an endothermic process, energy is absorbed
• The evidence that you see in the following
  examples is that the surroundings get cooler
• The net heat absorbed from the surroundings comes
  from the making and breaking of bonds during a
  chemical reaction (see p. 46-49).
• Evaporation
• Melting

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Water and Its Phase Changes