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Chemistry: the study of the properties and behavior of the Matter.
Matter: anything that has Mass and occupies Space.
Property: any characteristic that helps us recognize a particular type of matter
and to distinguish it from other types
Countless experiments have shown that the tremendous variety of matter is due
to combination of only about 100 very basic or elementary substances called
elements.
Chemistry relates the properties of matter to its composition, that is, to the
particular elements it contains.
Each element is composed of a unique kind of atom; atoms are the almost
infinitesimally small building block of matter.
The properties of matter relate to both the kinds of atoms the matter contains
(composition) and the arrangements of these atoms (structure).
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Atoms can combine to form molecules in which two or more atoms are joined
together in specific shapes:
Examples of molecules: water, ethanol, ammonia, oxygen, methane….
Even minor differences in the composition and structure of molecules can cause
profound differences in their properties.
Examples: ethanol and ethylene glycol.
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Classification of matter according to its composition
Matter
Pure substance
Mixtures
Air, gasoline,
sidewalk
Water, table salt
Compounds
Elements
Contain only one kind of atom
Contain two or more kinds of atoms
Water,
table salt
A pure substance is matter that has distinct properties and a composition that does not vary
from sample to sample; a pure substance, or simply a substance, can be either an Element or a
Compound.
Elements cannot be decomposed into simpler substances. On the molecular level, each
element is made of only one kind of atom.
Compounds are formed when elements interact with each other. For example when hydrogen
gas burns in oxygen gas, the elements hydrogen and oxygen combine to form the compound
water. So, compounds are substances composed of two or more elements; they contain two or
more kinds of atoms.
Mixtures are combinations of two or more substances in which each substance retains its own
chemical identity.
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Properties of matter
Physical properties: can be observed without changing the identity and
composition of the substance: color, odor, density, melting point,
boiling point and hardness.
Chemical properties: describe the way a substance may change or react
to form other substances: flammability, the ability of a substance to
burn in the presence of oxygen.
Physical and Chemical Changes
During a physical change, a substance changes its physical appearance
but not its composition: evaporation of water, change of state.
In a chemical change (chemical reaction) a substance is transformed
into a chemically different substance: hydrogen combining with oxygen
to form water.
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Separation of a mixture
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Atoms, Molecules and Ions
Atoms: the smallest units of matter that can undergo chemical
reactions.
The atom is composed mainly of 3 subatomic particles: protons,
electrons and neutrons.
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Atomic numbers, Mass Numbers, and Isotopes
What makes an atom of one element different from an atom of another element?
The atoms of each element have a characteristic number of protons; the number
of protons in the nucleus of each atom is called atomic number.
An atom has no net electrical charge; therefore # of protons = # of electrons.
All atoms of Carbon have 6 electrons and 6 protons. Carbon has atomic # of?
All atoms of Oxygen have 8 electrons and 8 protons. Oxygen has atomic # of?
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Atoms of a given element can have different # of neutrons; for example, most
atoms of carbon have 6 neutrons, but some have more and some have less.
Different isotopes of carbon are:
C
All atoms of an element have the same atomic #, but can have different mass #.
Atoms with identical atomic number but different atomic mass (that is, same
number of protons but different number of neutrons) are called isotopes.
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THE PERIODIC TABLE
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Molecules and Chemical Formulas
A molecule is an assembly of two or more atoms tightly bound together.
Many elements are found in molecular form; that is, two or more of the same type of atom are
bound together.
O2, H2, N2, halogens
Molecular Compounds are composed of molecules containing more than one type of atom.
H2O, CO2, CO, CH4
Most molecular compounds contain only nonmetals
Molecular formula of a compound indicates the numbers and types of atoms in a molecule.
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Picturing Molecules; the Structural Formula
The molecular formula of a substance summarizes the composition of the substance but does
not show how the atoms come together.
The Structural formula of a substance shows which atoms are attached to which within a
molecule.
In a molecule, the atoms are represented by their chemical symbols, and lines are used to
represent the bonds that hold atoms together.
A structural formula, however, does not depict the actual geometry of the molecule, that is, the
actual angels between the bonds.
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Ions and ionic compounds
The nucleus of an atom is unchanged by chemical processes.
Some atoms can readily gain or lose electrons to form charged particles called ions.
If electrons are removed from an atom, a positive ion, a cation, is formed. Example:
If an atom gains electrons, a negative ion, an anion, is formed. Example:
In general, metal atoms tend to lose electrons to form cations, whereas nonmetal
atoms tend to gain electrons to form anions.
Cations are Metal ions, whereas, anion are nonmetal ions.
In addition to simple ions, such as Na+ and Cl-, there are polyatomic ions, such as
NH4+, SO42-
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Predicting ionic charge
Atoms gain or lose electron to end up with the same number of electron as the noble gas
closest to them in the periodic table.
Noble gases are chemically very non-reactive and form very few compounds. This is because
their electron arrangements are very stable.
Nearby elements can obtain these same arrangements by losing and gaining electrons.
Ionic compounds
When electrons are transferred from one substance to another, ionic compounds are formed.
When elemental sodium reacts with elemental chlorine, an electron is transferred from a
neutral sodium atom to a neutral chlorine atom.
Na+ and Cl- ions are formed as a result of the electron transfer.
The ions of opposite charge attract and bind together to form an ionic compound: NaCl.
So, ionic compounds contain both + and – ions.
Ionic compounds are generally combinations of metals and nonmetals. In contrast, molecular
compounds are generally composed of nonmetals.
If we know the charges of ions forming an ionic compound, we can readily write the empirical
formula for an ionic compound.
Chemical compounds are always electrically neutral. So, the ions always occur in such a ratio
that total + charge equals total – charge.
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Examples: ionic compound formed from Mg and N
Naming Ionic compounds
1. Positive ions
a) Cations formed from metal atoms have the same name as the metal:
Na+, Zn2+, Al3+
b) If a metal can form different cations, the + charge is indicated by a Roman numeral in
parentheses following the name of the name of the metal:
Fe2+ iron (II)
Fe3+ iron (III)
Cu+
Cu2+
c) Cations formed from nonmetal atoms have names that end in –ium:
NH4+
H3O+
2. Negative Ions
a) If monoatomic, replace the ending of the name of the element with –ide
HO2N3A few simple polyatomic anions also have names ending in –ide
OHCNO22-
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Name and Formulas of Binary Molecular Compounds
1. The name of the element farther to the left in the periodic table is usually written first.
But, oxygen is always written last except when combined with fluorine.
2. If both elements are in the same group, the one having the higher atomic number is
named first.
3. The name of the second element is given an –ide ending.
4. Greek prefixes (mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca) are used to
indicate the number of atoms of each element.
Cl2O
N2O4
NF3
P4S10
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Nomenclature of organic compounds
Structural Isomers
Isomers have different structures with the same number of atoms.
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Chemical equations
We use chemical formula to write equations that represent chemical reactions.
Example of chemical reactions: striking a match, adding vinegar to a glass of water containing
baking soda, and caramelizing sugar. For all of these reactions we have visual evidence that
something has happened.
In this chapter, we are looking at some simple patterns of chemical reactivity.
Some simple chemical reactions are as follows: combination reaction, decomposition reaction,
and combustion reaction.
Chemical reactions are represented by chemical equations
2H2 + O2 → 2H2O
In any reaction, atoms are neither created nor destroyed. Therefore, we must have an equal
number of atoms of each element on each side of the arrow. When this condition is met, the
equation is balanced.
In balancing chemical equations, you should never try to change the subscript.
Balance the following chemical equations:
CH4 + O2 → CO2 + H2O
O2 + NO → NO2
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Some simple patters of chemical reactivity
How to predict the products of some chemical reactions knowing only their reactants?
Combination reactions:
A+B→C
Two or more substances react to form one product.
Example: elements combine to form compounds.
Mg (s) + O2 (g) → MgO (s)
This reaction is used to create bright flame used in fireworks.
Is the reaction balanced?
What kind of a compound is MgO?
Write the products of the following reactions:
C (s) + O2 (g) →
N2 (g) + H2 (g) →
CaO (s) + H2O (l) →
Decomposition Reaction:
One substance undergoes a reaction to produce two or more substances:
CaCO3 (s) → CaO (s) + CO2 (g)
C→A+B
KClO3 (s) → KCl + O2
PbCO3 →
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Ca(OH)2 →
Predict the product(s) of the following reactions:
a) Combination of lithium metal with fluorine gas
b) When solid barium carbonate is heated
Combustion in Air:
Combustion reactions are rapid reactions that produce a flame.
Examples: burning of hydrocarbons combine with O2 from air to form CO2 and H2O.
C3H8 (g) + O2 (g) → CO2 (g) + H2O (g)
Write the equation of combustion reaction of methanol.
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Formula and Molecular Weights
The formula weight of a substance is the sum of atomic weights of each atom in its
chemical formula.
F.W. of H2SO4 = 2(A.W. of H) + (A.W. of S) +4(A.W. of O)
= 98.1 amu
F.W. of Na is
M.W. of C6H12O6 is
F.W. of NaCl is
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Avagadro’s Number and the Mole
Even the smallest samples we deal with have enormous numbers of atoms, ions or molecules.
5 mL of water contains 2×1023 water molecules.
So, chemists have devised a special counting unit for describing such large numbers.
A mole is the amount of the matter that contains in exactly 12 grams of isotopically pure 12C.
From experiments, this number is determined to be 6.02 × 1023
1 mol 12C = 6.02 × 1023
12C
atoms
1 mol H2O molecules =
1 mol NO3- ion =
Calculate the number of H atoms in 0.35 mol of C6H12O6
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Molar Mass
Mass in grams of one mole of a substance; the mass in grams per mol; g/mol.
Just like a dozen eggs has different weight than a dozen elephants.
1 atom of 12C has a mass of 12 amu ; 1 mol of 12C has a mass of 12 grams.
1 atom of Cl has an atomic mass of 35.5 amu; 1 mol of Cl has a mass of 35.5 g.
The molar mass (g/mol) of any substance is always numerically equal to its formula weight (in
amu).
NaCl has a formula weight of 58.5 amu and a molar mass of 58.5 g/mol.
BaCl2 has FW and MW of
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Chap 3
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2C6H14O4(l) + 15O2(g) → 12CO2(g) + 14H2O(l)
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Preparation of 0.250 M K2CrO4(aq) - Example 4-9 illustrated
Visualizing the dilution of a solution
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Preparing a solution by dilution - Example 4-10 illustrated
(a) A pipet is used to withdraw a 10.0-mL sample of 0.250 M K2CrO4(aq). (b) The pipetful of
0.250 M K2CrO4(aq) is discharged into a 250.0-mL volumetric flask. (c) Water is then added
to bring the level of the solution to the calibration mark on the neck of the flask. At this
point, the solution is 0.0100 M K2CrO4(aq).
General Properties of Aqueous Solutions
Substances dissolve in water and can exist in as ions, molecules or a mixture of the two.
A solution in which water is the dissolving medium is called an aqueous solution.
Water is the medium for most of the chemical reactions that take place within us or around us.
Limestone caves are formed by the dissolving action of underground water on CaCO3 in the
limestone.
CaCO3 (s) + H2O (l) + CO2 (g) → Ca(HCO3)2 (aq)
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General properties of Aqueous Solutions
Solutions are homogeneous mixtures of two or more substances.
The substance present in the greatest quantity is called Solvent.
The other substances are called Solutes.
Electrolytic Properties
The ability of a solution to conduct electricity depends on the number of ions it contains.
An electrolyte solution contains ions that serve as charge carriers.
The conductivity of NaCl solution indicates the presence of ions in the solution. The solution
contains Na+ and Cl-. Each ion is surrounded by water molecules.
The lack of conductivity of sucrose solution indicates the absence of ions. The solution contains
neutral molecules C12H22O11of surrounded by H2O molecules.
A substance whose aqueous solution contains ions is called an electrolyte.
A substance whose aqueous solution does not contain ions is called a nonelectrolyte.
Ionic solids dissolve in water and dissociate into their component ions. These ions become
surrounded by water molecules. The ions are said to be solvated. The solvation process helps
stabilize the ions in the solution and prevent them from recombining.
We can predict the nature of the ions present in a solution of an ionic compound simply from
the name of the compound:
Sodium sulfate
Calcium chloride
Molecular compounds in water
When molecular compounds dissolve in water, the solution usually consists of intact molecules.
So, most molecular compounds are nonelectrolytes.
Examples: CH3OH, table sugar.
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In (a), there are no ions present to speak of-only molecules. Methanol (methyl alcohol), is a
nonelectrolyte in aqueous solutions. In (b), the solute is present almost entirely as individual
ions is a strong electrolyte in aqueous solutions. In (c), although most of the solute is present as
molecules, a small fraction of the molecules ionize. is a weak electrolyte in aqueous solution.
The molecules that ionize produce acetate ions and ions, and the ions attach themselves to
water molecules to form hydronium ions,
However, Acids, dissolve in water and dissociate into ions:
HCl (aq) → H+ (aq) + Cl- (aq)
Notice the single arrow represents 100% ionization.
Strong and weak electrolytes
Strong electrolytes are those solutes that exist in solution completely as ions. Examples are
soluble ionic compounds and strong acids like HCl
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Refer to Table 4.1 solubility guidelines for common ionic compounds p125
Weak electrolytes are those solutes that exist in solution mostly in the form of molecules with
only a small fraction in the form of ions. Examples are weak acids like CH3COOH; only about 1%
of CH3COOH is present as H+ (aq) and CH3COO- (aq) ions.
When a weak electrolyte ionizes in water, we write the reaction in the following manner:
CH3COOH (aq)
CH3COO- (aq) + H+ (aq)
The half arrows in both directions mean that the reaction is significant in both directions.
Acids and Bases are also common electrolytes.
Acids
Acids are substances that ionize in aqueous solutions to form H+
Examples of strong acids:
HNO3
HCl
H2SO4
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Weak acids only partially ionize in aqueous solutions
CH3COOH
Bases
Bases are substances that accept H+ or react with H+.
Bases produce OH- ions in water.
Ionic hydroxides, NaOH, Ca(OH)2, KOH, are among the most common bases. Most common
strong bases are the hydroxides of the alkali metals, group 1A, and hydroxides of Ca2+, Sr2+
and Ba2+.
Compounds that do not have OH can also be bases. Like ammonia NH3
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NH3 (aq) + H2O (l)
Ammonia is a weak electrolyte as only a small fraction (about 1%) of NH3 ionizes to form NH4+
and OH-.
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Neutralization reactions and Salts
When a solution of an acid and a solution of a base are mixed, a neutralization reaction is
occurs. The products of this reaction are salt and water.
In acid-base reactions H+ and OH- come together to form H2O.
HCl (aq) + NaOH (aq) →
Mg(OH)2 (aq) + HCl →
Carbonates and bicarbonates react with acids to produce CO2
HCl (aq) + NaHCO3 (aq) →
Na2CO3 (aq) + CH3COOH (aq) →
Predict what gas is produced when HCl reacts with a metal sulfide such as Na2S?
Oxidation Numbers
The concept of oxidation number helps us keep track of the electrons gained by the substance
reduced and electrons lost by the substance oxidized in a redox reaction.
Each atom in a neutral molecule or charged species is assigned an oxidation number.
Here are the rules to assign oxidation numbers or oxidation states:
1. For an atom in its elemental form, the oxidation number is always zero.
H in H2, P in P4 has an oxidation number of zero.
2. For ant monatomic ion the oxidation number equals the charge on the atom.
K+ has the oxidation number of +1, Ca2+ has the oxidation number of +2, S2- has the
oxidation number of -2.
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3. Nonmetals usually have –ve oxidation number.
a) The oxidation number of oxygen is usually -2 in both ionic and molecular
compounds.
Exception: in peroxides, compounds containing O22-, each O has oxidation number of
-1.
b) The oxidation number of H is usually +1 when bonded to nonmetals and -1 when
bonded to metals.
c) The oxidation number of F is -1 in all compounds; other halogens have an oxidation
number of -1 in most binary compounds. When combined with O, as in oxyanions,
however, they have +ve oxidation number.
4. The sum of the oxidation numbers of all atoms in a neutral compound is zero.
The sum of the oxidation numbers in a polyatomic ion equals the charge of the ion.
In H3O+ …..
Determine the oxidation number of S in the following:
SCl2, Na2SO3, SO42-
Oxidation- Reduction Reactions
In redox reactions electrons are transferred between reactants.
Example: corrosion (rusting) of iron, that is, the reaction of iron with oxygen in the presence of
water.
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What we call corrosion is the reaction between a metal and some substances in the
environment.
Calcium is vigorously attacked by acids to form Ca2+:
Ca (s) + H+ →
Ca becomes more +ve, loses electrons therefore it is Oxidized.
When a metal corrodes, it loses electrons and becomes more +vely charged.
Many metals react directly with O2 from air to form metal oxides.
Ca (s) + O2 (g) →
In the reaction above Ca has lost electrons and is more +ve, while O2 had gained electrons to
become more –ve.
When an atom or molecule gains electrons and becomes more –ve, it is reduced.
So, the loss of electron by a substance is called oxidation. The gain of electrons by a substance
is called reduction.
A reactant that loses electrons is oxidized.
A reactant that gains electrons is reduced.
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Example of displacement reactions in which the ion is displaced or replaced through oxidation
of an element.
Many metals react with acids to produce salt and H2 gas.
Mg (s) + HCl (aq) →
Metals can also be oxidized by aqueous solutions of various salts.
Molecular equation: Fe (s) + Ni(NO3)2 (aq)
Net ionic equation:
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Earth's crust
Other
9%
Silicon
26%
Aluminum
8%
Oxygen
49%
Iron
5%
Calcium
3%
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Patterns of chemical reactivity
The products of a reaction can sometimes be predicted by using the periodic table and from
knowledge of similar reactions
1. Using the periodic table
Na + H2O
K is in the same group as Na; so, they are expected to behave similarly
2. Combustion in air: involve reaction of a hydrocarbon such as CH4, C2H6 (contain only C
and H) with oxygen to form H2O and CO2
3. Combustion of compounds containing C, H and O also produce CO2 and H2O
Exercise: write the balanced chemical equation for the reaction that occurs when methanol
CH3OH is burned in air
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4. Decomposition Reactions: in which a single reactant forms two or more products usually
when heated.
Decomposition of metal carbonates
Examples:
Write the chemical equations for the following reactions:
a. Combination when Li metal and fluorine F2 gas react.
b. Decomposition reaction when solid barium carbonate is heated.
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