Unit 1: Chemistry P' 134279

1
Unit 1 Chemistry (P. 134-279)

• Patterns and Compounds
• Periodic Table, Naming, Balancing Equations
• Chemical Reactions
• Energy, 4 Types, Combustion
• Acids and Bases
• Properties, pH, Reactions
• Chemical Reactions in the Environment
• Factors affecting Rates, Chemicals and Us

2
Classification of Matter
3
Classification of Matter

• PURE SUBSTANCEA substance with a fixed
  composition and constant properties
• ELEMENT A substance that cannot be broken down
  into simpler substances by chemical means.
  Atoms are the simplest particles that cannot be
  broken down by chemical means
• COMPOUND A substance that is made up of two or
  more different atoms (molecules). These
  substances can be broken down only by chemical
  means.
• MIXTURE A mixture consists of two or more kinds
  of matter, each keeping its own characteristic
  properties.
• SOLUTIONS A mixture that is homogeneous. If the
  solution is a liquid or gas, it is transparent.
• MECHANICAL MIXTURE A heterogeneous mixture with
  parts that are visibly distinguishable

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Metals Metalloids Nonmetals
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Metals vs. Non Metals

• Shiny
• Ductile
• Malleable
• Conducts Electricity
• Conducts Heat

• Dull
• Brittle
• Does Not Conduct Electricity
• Does Not Conduct Heat

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Elements and the Periodic TableA Review

• Every Element has a unique
• Name
• Symbol
• Atomic Mass Number (A)
• Represents the number of protons number of
  neutrons
• Atomic Number (Z)
• Represents the number of protons and the number
  of electrons in a neutral atom
• The number of neutrons can be calculated by
  subtracting A - Z

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Atomic Number Atomic Mass
Number of electrons, protons
2
Atomic Number
He
4
Number of Protons and Neutrons
Atomic Mass
Number of Neutrons Atomic Mass Atomic Number
8
Example
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Grouping of Elements

• Elements are subdivided into
• groups or "families" (vertical columns)
• and periods (horizontal rows)
• Metals elements are on the left
• Non-metal elements on the right
• separated by a dark "staircase line".
• Elements bordering this division line exhibit
  some properties of both metals and non-metals and
  are called metalloids.
• Copy table 5.1 on page 140 into your notes.

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Group Names
Alkali Metals
Alkaline Earth Metals
Group 3 Carbon Group Nitrogen Group Oxygen Group
Halogens
Noble Gases
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Bohr-Rutherford Diagrams

• The following information is required
• 1. Number of Electrons
• Is the same as the number of protons in a neutral
  atom. The electrons are organized into shells in
  the following order.
• up to 2 electrons in the first shell
• up to 8 electrons in the second shell
• up to 18 electrons in the third shell
• up to 32 electrons in the fourth shell
• 2. Number of Protons
• Is the same number as the atomic number
• 3. Number of Neutrons
• Can be determined by subtracting the atomic
  mass from the atomic number

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Bohr Rutherford Diagram
Nucleus with protons and Neutrons
18p 36n
Electron Orbits (shells) with a 2,8,8, pattern
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Predicting Chemical Reactivity

• Elements with 8 electrons in their outer energy
  level appear to have a special significance.
  Elements with this arrangement do not react
  easily and are considered stable.
  
  
• All noble gases (Neon, Krypton, Xenon, Radon)
  have 8 electrons in their outer energy level and
  are very non-reactive elements (Helium is a
  special gas that is very stable with 2 electrons
  in its first level).
• All elements want to be stable and therefore want
  to gain or lose electrons in order to achieve a
  stable 8 configuration (Stable Octet).
• Whenever an atom gains or loses electrons they
  become negative or positive and they are called
  ions.

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Two main factors determine chemical
activity(reactivity)

• 1) The number of electrons in the outer energy
  level
• i) Elements with 1-3 electrons on outer level
  lose electrons (become positive)
• ii) Elements with 5-7 electrons in outer level
  gain electrons (become negative)
• iii) Elements with 4 electrons in outer level are
  special (tend to become positive)
• 2) The number of energy levels
• As the number of energy levels increase, the
  attraction between those electrons in the
  outermost energy level and the positive nucleus
  decrease

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Ions To gain or lose an Electron

• Positively Charged Cations
• When a neutral atom gives up one or more
  electrons, the positively charged ion that
  results is called a Cation.
• For example
• Negatively Charged Anions
• When a neutral atom gains one or more electrons,
  the negatively charged ion that results is called
  an Anion.
• For example

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Electron Dot Diagrams

• A Bohr-Rutherford diagram represents an atom and
  all its electrons.
• A simpler way to represent atoms and ions of atom
  is with electron dot diagrams
• Electron Dot Diagrams show only the outer energy
  level (valence shell) of an atom. Only these
  electrons are represented because they are
  responsible for an atoms chemical properties.
  For example

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Lewis Dot / Electron Dot diagrams
N
C
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Chemical Bonds Forming Compounds

• Most substances on earth do not exist as
  elements, they are composed of two or more
  different elements joined together to make
  compounds.
• When two atoms collide, valence electrons on each
  atom interact. A chemical bond forms between
  them if the new arrangement of their valence
  electrons have less energy than their previous
  arrangement.
• For many atoms that new arrangement of their
  electrons will be that of their closest noble
  gas.
• Atoms may acquire a valence shell like that of
  its closest noble gas in one of three ways
• 1. An atom may give up electrons and forma ion
• 2. An atom may gain electrons and form an ion
• 3. An atom may share electrons

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Ionic Compounds

• Substances held together by ionic bonds are
  called
• Ionic compounds e.g. NaCl, KCl. Ionic Bonds
• occur because of the attraction of cations and
  anion
• for each other. Electrons are transferred between
  the
• atoms during bond formation.
• Properties include
• High melting point (i.e. strong bonds)
• Conduct electricity when dissolved in water
  or molten
• Form crystal lattice structures
• Soluble in water

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Molecular Compounds

• Substances that are composed of molecules are
  called molecular compounds. Many non-metals form
  compounds with other non-metals. When this
  occurs there is no transfer of electrons between
  the two atoms instead they share electrons
  forming a covalent bond.
• Although bond between atoms are strong, bonds
  between molecules are weak. eg. Moth crystals,
  nitrogen gas etc.
• Properties Include
• Low melting and boiling points
• Often have an odour
• Dont conduct heat
• Dont conduct electricity (non-electrolytes)
• Diatomic molecules (e.g. O2, F2 etc.) are also
  the
• result of covalent bonds.

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Chemical Naming and Formulas

• Binary Ionic
• Transition Metals
• Stock versus Classical
• Polyatomic Ions
• Binary Molecular

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General Rules

• The Metal is always written first
• The nonmetal suffix in a compound is either ide
  or ate
• Every compound must be electrically neutral
• All Positive charges must equal Negative charges

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Binary Compounds Formula to Name

• Composed of two Elements
• One metal and one nonmetal
• Write the name of the metal first unchanged
• Write the name of the nonmetal second
• Change the ending to an ide
• LiCl ? Lithium Chloride
• MgI2 ? Magnesium iodide

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Binary Compounds Name to Formula

• Write the symbol for each element with the metal
  written first
• Find the ionic charge for each element
• Cross the number value of the charge and place it
  as the subscript of the other element
• Reduce the values to lowest ratio
• Magnesium Oxide ? Mg2 O2-
• Mg2O2
• MgO

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Transition Metals Groups 3-12Name to formula

• Almost all are able to form more than one cation
• When writing the formula the charge of the metal
  cation will be indicated by roman numerals after
  the metal
• Lead (III) chloride ? PbCl3
• Iron (II) oxide ? FeO

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Transition Metals Groups 3-12Formula to Name

• Finding the charge on the metal can be done two
  ways
• Reverse Cross-Over Method
• The subscript of the nonmetal becomes the charge
  of the metal
• Sometimes the charge is misleading
• Charge Balancing
• Charge Subcript of the nonmetal multiplied by
  the charge of the nonmetal divided by the
  subscript of the metal

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Chemical Equations and Reactions

• A chemical equation is a description of a
  chemical
• reaction using chemical symbols, not words
• Steps
• 1) The reactants are written first
• 2) The products are written second
• 3) The state for each atom is indicated
• (g) gas, (s) solid, (l) liquid, (aq) aqueous
  
• 4) The reactants and products are separated by
  an
• "arrow" ( ? )
• e.g. Word Equation
• Hydrogen gas plus chlorine gas produces
• hydrogen chlorine gas
• e.g. Chemical Equation
• H2(g) Cl2(g) ? HCl(g)

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Balanced and Unbalanced Chemical Equations

• The Law of Conservation of Mass states
• Matter cannot be created or destroyed it can
• only be changed from one form to another.
• Therefore, the number of atoms in the reactants
• must equal the number of atoms in the products
• An unbalanced or skeleton equation does not
• follow the Law of Conservation of Mass. The
• number of atoms on the left side (reactants)
• does not equal the atoms on the right side
• (products)
• e.g. H2(g) Cl2(g) ? HCl(g)
• 4 atoms (2 H, 2 Cl) 2 atoms(1 H,
  1 Cl)

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• A balanced chemical equation follows the Law
• of Conservation of Mass. The number of atoms
• on the left side (reactants) equals the atoms on
• the right side (products)
• e.g. 1H2(g) 1Cl2(g) ? 2HCl(g)
• 4 atoms (2 H, 2 Cl) 4
  atoms(2 H, 2 Cl)

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Writing Balanced Chemical Equations

• 1. Write the chemical formula for each reactant
  and product followed by the state of each solid
  (s) liquid (l) gas (g) aqueous(aq)
• 2. Adjust the numbers of molecules until there
  are the same number of atoms of each type on both
  sides of the equation. This balances the mass of
  both the reactants and products.
• 3. Usually, balancing is easiest when hydrogen
  and oxygen atoms are left until the end
• NOTE
• Do not change the subscript in a formula to
  balance an
• equation. Changing these numbers changes the
• molecular structure of the molecule.

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Energy Changes and Chemical Reactions

• Chemical reactions, physical changes of state and
  
• dissolving processes often involve energy
  changes.
• Exothermic Processes
• Processes that release energy (e.g. heat and
• light) and increase the temperature of the
• surroundings.
• Endothermic Processes
• Processes that absorb energy and decrease the
• temperature of the surroundings.

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Factors Affecting Chemical Reaction Rate

• The Rate of Reaction is defined as
• The time it takes for a given product to form, or
  for
• given amounts of reactant to react.
• Reaction rate is determined by
• i. Measuring how fast reactants are used up.
• ii. Measuring how fast the products are formed.
• Factors affecting Reaction Rate
• 1. Concentration and Reaction Rate
• ? Concentration (amount of substance in a given
  volume) ? Rate
• 2. Surface Area and Reaction Rate
• ? Surface Area (area exposed) ? Rate

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• 4. Catalysts and Reaction Rates
• A Catalyst is defined as
• A substance that speeds up the rate of a chemical
  
• reaction without being used up in the reaction.
• Catalyst lower the energy required to break the
• bonds that hold substances together. Examples
• include enzymes (biological catalysts),
• platinum, rhodium and palladium (chemical
• catalyst used in catalytic converters)

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Types of Chemical Reactions

• There are four basic patterns that most chemical
• reactions follow
• 1) Synthesis Reactions
• This type of reaction fits the general pattern
• A B ? AB
• e.g. N2(g) 3H2(g) ? 2NH3(g)
• CaO(s) H2O(l) ? Ca(OH)2
• A synthesis reaction involves the formation of a
• new compound from simpler elements or
• compounds
• Combustion reactions (involving the reaction
• with O2) are examples of Synthesis Reactions

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2) Decomposition Reactions

• These type of reactions are opposite to direct
• combinations. They fit the general pattern
• AB ? A B
• e.g. CuCO3(s) ? CuO(s) CO2(g)
• 2KClO(s) ? 2KCl(s) 3O2(g)
• A decomposition reaction involves the breaking
• down of a compound into simpler compounds or
• elements

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3) Single Displacement Reactions

• A single displacement or substitution reaction
• fits the general pattern of
• A BC ? AC B
• This type of reaction involves a change in
• partners. One element displaces or knocks off
• another element in a compound..
• e.g. Zn(s) 2HCl(aq) ? ZnCl2(aq) H2(g)
• 3C(s) Fe2O3(s) ? 3CO(g)
  2Fe(s)

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4) Double Displacement Reactions

• A double displacement reaction fits the following
  
• general pattern
• AB CD ? AD CB
• This type of reaction involves a change of both
• partners. The cation (positive element or
  polyatomic
• ion) of one compound changes place with the
  cation
• of the second compound.
• e.g. Na2S(aq) ZnCl2 (aq) ? ZnS(s)
  2NaCl(aq)
• AgNO3(aq) KBr(aq) ? AgBr(s)
  KNO3(aq)
• SF4(s) 2H2O(l) ? SO2(g)
  4HF(aq)

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• Carbon Chemistry
• Organic Chemistry The study of carbon
• containing compounds and their properties e.g.
• hydrocarbons
• When hydrocarbons (contain carbon and hydrogen)
• are burned in enough oxygen complete combustion
• occurs.
• Hydrocarbon oxygen gas ? carbon dioxide
  water E
• (good supply)
• If hydrocarbons are burned in a poor supply of
• oxygen, incomplete combustion occurs.
• Hydrocarbon oxygen gas ? carbon dioxide
  water E
• (poor supply)
  carbon monoxide residue

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Classification of Substances by Their Behaviour

• The process of grouping substances
• according to common properties is called
• classification.
• Previously we have classified substances
• according to
• i) State (e.g. solid, liquid or gas)
• ii) Composition (e.g. pure substances, mixtures
  etc.)
• Matter can also be classified by
• chemical behaviour.
• Acids and bases make up two classes of
• compounds that have been classified by
• their chemical behaviour.

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Acids and Bases

• Acids
• An acid is a compound that dissolves in water to
• produce hydrogen ions (H ) in solution. e.g.
  HCl
• Bases
• A base is a compound that dissolves in water to
• produce hydroxide ions in solution (OH -) e.g.
  NaOH
• Copy Table 7.3 Acids and Bases A Summary
• found on page 230 in your text.

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Preparation of Common Acids

• A common way to prepare an acid is to react a
  nonmetal oxide
• with water. An oxide is an element combined with
  only
• oxygen e.g.
• sulphur trioxide water ? sulphuric acid
• carbon dioxide water ? carbonic acid
• Some common acids in the laboratory include
• i) sulfuric acid ( H2S04 )
• ii) nitric acid (HNO3)
• iii) hydrochloric acid (HCl)
• iv) acetic acid, (CH3COOH)
• Other common acids include

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Preparation of Common Bases

• A common way to prepare a base is to react a
  metal oxide
• with water. e.g.
• sodium oxide water ? sodium hydroxide
• calcium oxide water ? calcium hydroxide
• Some common bases in the laboratory include
• i) Sodium hydroxide (NaOH)
• ii) Calcium hydroxide (Ca(OH)2)
• iii) Potassium hydroxide (KOH)
• iv) Magnesium hydroxide (Mg(OH)2)

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Indicators

• An indicator is a chemical that changes colour as
  
• the concentration of H (aq) and OH- (aq)
• changes. e.g.
• i) Litmus
• blue litmus turns red in acid
• red litmus turns blue in base
• ii) Phenolphthalein
• turns pink in base
• Indicators can be made from flowers, fruits,
• vegetables, leaves (e.g red cabbage, tea etc.)
• Synthetic Indicators are more easy to use than
• natural indicators because they
• last longer than natural indicators
• can be produced in large quantities
• e.g. bromothymol blue (BTB)
• phenolphthalein

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The pH Scale

• The pH scale describes the "strength of the
• hydrogen ion (H)".
• The scale is numbered from 0 to 14
• acids have a pH less than 7 H
  gt OH-
• bases have a pH more than 7 H lt
  OH-
• neutral substances have a pH of 7 H
  OH-
• The change in 1 pH unit represents a tenfold
• increase in the concentration of hydrogen ions
  in
• solution. e.g.
• A pH of 2 is 10 x's stronger than a pH of 3
• A pH of 2 is stronger than a pH of 5
• A pH of 2 is _ stronger than a pH of 7
• pH can be estimated using pH paper or measured
• using a pH meter (measures electric properties)

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The Strength Of Acids And Bases

• The strength of an acid or base is dependant on
  two
• factors
• 1. Concentration
• The concentration of an acid or base is the
  amount
• of the pure substance dissolved in 1 L of water.
• 2. Ionization
• When acids and bases are dissolved in water, they
  
• ionize (break apart into charged particles). The
  term
• Percent Ionization refers to the number of
• molecules that will ionize for every 100
  molecules
• that dissolve. e.g. HCl H2O ? H3O Cl-
• Solutions that form ions in water are called
• electrolytes. Electrolytes conduct electricity.

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• The Strength of Acids
• Strong acids ionize completely in water e.g
  H2SO4
• Weak acids ionize partially in water e.g.
  CH3COOH
• The Strength of Bases
• Strong Bases ionize completely in water e.g
  NaOH
• Weak Bases ionize partially in water e.g. NH3

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Neutralization

• Neutralization occurs when hydroxide ions
• (base) and hydrogen ions (acid) are mixed to
• make water. The general word equation is
• Acid Base ? Water Salt
• e.g
• hydrochloric sodium ? water
  sodium chloride
• acid hydroxide
• (HCl) (aq) ( NaOH)(aq) ? (
  H2O)(l) ( NaCl) (aq)
• After neutralization, the solution no longer has
• a high concentration of either ion.

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Soaps and Detergents

• What makes up soap ?
• 1. fatty acid (lipid)
• 2. strong base (NaOH)
• The word equation is
• fat base ? soap glycerol
• Soap curds cling as scum to whatever it comes
• into contact with, and does not rinse away
  easily.
• This problem led to the development of synthetic
  
• detergents called syndets. Advantages include
• 1. good at removing dirt
• 2. more soluble in water
• 3. prevented dirt from collecting back onto
  clothes
• 4. did not form a curd
• 5. mild to hands and fine fabrics
• 6. less expensive (made from plant oils and
  animal fats)

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• How soap cleans
• A soap or detergent molecule consists of
• two ends
• 1. Hydrophillic (water loving)
• The end with the sodium ion is attracted to
  water and becomes soluble
• 2. Hydrophobic (water hating)
• Hydrocarbon end is attracted to insoluble
  dirt (grease)
• on clothes etc.
• For example