This document provides an introduction to inorganic chemistry concepts including definitions of matter, mass, elements, atoms, and the periodic table. Key topics covered include common elements, atomic structure, the periodic table, chemical bonding including covalent, ionic, and hydrogen bonds. Examples are given throughout to illustrate important inorganic chemistry concepts and chemical bonding.
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ElementElement
• A substance that cannot be broken down chemically into
simpler substances
• About 92 naturally occurring elements
– Example – gold, magnesium, neon, etc.
• About 25 elements necessary for life
• Six most common elements
– Carbon, hydrogen, oxygen, nitrogen . . . calcium and
phosphorus
– >99% of living things
• Trace elements – necessary in only small (i.e. trace)
amounts
– Example – Iodine
• Necessary for proper thyroid function
• Goiter caused by iodine deficiency
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AtomAtom
• Simplest particle of an element that retains
the properties of that element
• A given atom is unique to a given element
• Atoms composed of smaller particles
called subatomic particles
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Compounds vs. MoleculesCompounds vs. Molecules
• Compound - A substance composed of
two or more elements.
– Ex. NaCl – sodium chloride (table salt)
• Molecule – A substance composed of two
or more atoms, can be the same element.
– Ex. O2 – oxygen (diatomic)
– However, sodium chloride is also a molecule
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Important DefinitionsImportant Definitions
• Atomic Number
– # of protons
– Unique for each element
• Mass number
– Sum of the protons and
neutrons
• Atomic Mass
– Average of masses of all of
the isotopes of an element
– Usually a decimal number
very close to the mass
number
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Periodic TablePeriodic Table
• Contains horizontal rows
called periods
– Ascending atomic # from
left to right.
• Contains vertical columns
referred to as groups.
– We are concerned with
groups IA-VIIIA
– Elements in a group have
similar chemical and
physical properties
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Group IA – Alkali MetalsGroup IA – Alkali Metals
• Strong metallic
qualities
• Highly reactive
– Not found alone in
nature
• One valence electron
• Tendency to lose one
e- when reacting.
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Group IIA – Alkaline EarthGroup IIA – Alkaline Earth
MetalsMetals
• Strong metallic
qualities
• Very reactive
– Not found free in
nature
• Harder and denser
than alkali metals
• Two valence e-
• Tendency to lose 2 e-
during reactions
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Group VIA - ChalcogensGroup VIA - Chalcogens
• More varied in properties
– Oxygen, Sulfer –
nonmetals
– Selenium, tellurium –
metalloids
– Polonium – metal
• Very Reactive
• Six valence e-
• Tendency to gain two
electrons
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Group VIIA - HalogensGroup VIIA - Halogens
• All nonmetals
• Often found in
diatomic state (F2)
• Very reactive
• Seven valence e-
• Tendency to gain 1 e-
during reactions
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Metals vs. NonmetalsMetals vs. Nonmetals
• Metals
– Solid at room temperature
– Conduct heat and electricity well
– Malleable (sheets)
– Ductile (wires)
– Lustrous (shiny)
– High melting/boiling points
• Nonmetals
– Opposite of metals
• Metalloids
– Some qualities of both metals and nonmetals
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Did you know?Did you know?
• Atoms are mostly empty space?
– If the nucleus of an atom was the size of a golf
ball, the nearest electron would be roughly 1 km
away!
• The nucleus of an atom is extremely dense.
– The same size nucleus would have a mass of
approximately 2.5 billion tons!
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Some More Things to KnowSome More Things to Know
• All atoms of a given element have the
same # of protons
• All atoms are considered neutral in charge
unless designated with a symbol of
charge, in which case they are considered
an ion; # electrons = # protons except in
ions
• The # of neutrons is equal to or greater
than the number of protons.
– mass # - atomic # = # neutrons
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3 Major Subatomic Particles3 Major Subatomic Particles
• Proton
– Positive charge
– 1 x 10 -24
grams (about 1 dalton)
– Located in the nucleus
• Neutron
– Neutral (no charge)
– 1 x 10 -24
grams (about 1 dalton)
– Located in the nucleus
• Electron
– Negative charge
– 1/2000 the mass of a proton or neutron
– Moving in orbitals around the nucleus at about the speed of light
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Energy Levels (electron shells)Energy Levels (electron shells)
• Electrons exist at varying energy levels
• The further they are from the nucleus, the
more energy they have
– Think centripetal force
• Electrons tend to occupy the lowest
energy level (closest to nucleus) possible
• Electrons can be “excited” to higher
energy levels for very brief periods
– Example: Light energy during photosynthesis
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Electron Configuration and ChemicalElectron Configuration and Chemical
PropertiesProperties
Why atoms reactWhy atoms react
• It’s all about the # of valence electrons!
• Valence electron shell is the outermost
shell (that contains electrons)
• A full valence shell = inert (stable electron
configuration)
• Anything else = reactive (unstable electron
configuration)
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Here’s the DealHere’s the Deal
• 1st
energy level
– Full (stable) with 2 electrons
• 2nd
energy level
– Full (stable) with 8 electrons
• 3rd
energy level
– Full (stable) with 8 electrons
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Isotopes
• Most elements have at least 2 isotopes,
some have several.
• Isotopes vary in the # of neutrons only.
• Example: Carbon has 3 isotopes
– 12
C – stable (6 neutrons)
– 13
C – stable (7 neutrons)
– 14
C – radioactive (8 neutrons)
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Uses of Radioactive IsotopesUses of Radioactive Isotopes
• Dating fossils
– Carbon – 14
• Measure half-life (5730 years)
• Medical tracers
– Iodine – 131
• Various types of sensors can detect radiation.
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Answers
• Lithium generally has
– 3 protons
– 3 electrons
– 4 neutrons
• Oxygen generally has
– 8 protons
– 8 electrons
– 8 neutrons
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Lewis Electron Dot DiagramsLewis Electron Dot Diagrams
• Show the electron configuration for only
the valence e- for an atom
• Steps
– Write the symbol of the atom
– Make a dot for each valence e- (use the “four
sides” of the symbol)
– Only one rule – don’t pair up e- until after all
four orbitals have one e- each
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4 Major Types of Bonds4 Major Types of Bonds
• Strongest to weakest
– Covalent bonds
– Ionic bonds
– Hydrogen bonds
– van der Waals interactions
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Covalent BondsCovalent Bonds
• Strongest
• Generally occurs when two nonmetals interact
• A pair, or pairs, of e- are shared
• Single covalent bond
– One pair of e- shared between two atoms
– Represented by a single line in structural formula
• Double covalent bond
– Two pairs of e- shared between two atoms
– Represented by a double line in structural formula
• Triple covalent bond
– Three pairs of e- shared between two atoms
– Represented by a triple line in structural formula
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Polar vs. Nonpolar Covalent BondsPolar vs. Nonpolar Covalent Bonds
• It’s all about electronegativity
– Electronegativity
• The affinity an atom has for electrons
– i.e. How strongly it pulls on both its own e- and the e- of other
atoms
• All atoms are electronegative, some more than
others
• Polarity, whether or not a molecule is polar or
nonpolar, can have a big effect on the behavior
of the molecule.
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Nonpolar Covalent BondsNonpolar Covalent Bonds
• Occurs between two atoms of the same
electronegativity.
• Electrons are shared equally
– Both atoms are pulling with the same force
• Examples – eneg = electronegativity
– O=O (O2)
• Same atom – same eneg
– C—H
• Carbon and hydrogen have the same eneg
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Polar Covalent BondsPolar Covalent Bonds
• Occurs between two atoms of differing eneg
• Electrons are not shared equally
– i.e. e- spend more time around one atom than the
other
• This creates a slight polarity of charge in the
molecule
– More eneg atom gains slightly negative charge
– Less eneg atom gains a slightly positive charge
• Note – oxygen is the big one here
• Example
– H – O bond
– Hydrogen is slightly positive, oxygen slightly negative
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Ionic BondsIonic Bonds
• Also strong
– Relatively weak around water
• Around water, ionically bonded substances dissociate into
ions
• Generally occur between a metal and a
nonmetal
– Metal loses electron, nonmetal gains electron
• Electrons are not shared, they are transferred
from one atom to another
• Differences in eneg are great
• Ions (charged particles) are formed
• An ionic bond is an attraction between
oppositely charged ions.
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Examples of Ionic BondsExamples of Ionic Bonds
• Na + Cl Na+
+ Cl-
NaCl
– Cl steals an e- from Na, gains a 1- charge and leaves
Na with a 1+ charge. The oppositely charged ions are
attracted.
• Mg + 2F Mg2+
+ 2F-
MgF2
– Two fluorinessteal 1 e- each from Mg, gain a 1-
charge and leave Mg with a 2+ charge. The
oppositely charged ions are attracted.
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Hydrogen BondsHydrogen Bonds
• Hbonding is an attraction between the slightly
positively charged atom in one polar bond and
the slightly negatively charged atom in a
different polar bond
• Occur only between polar molecules or polar
regions of molecules
• Weak, short-lived bonds (still very important)
• This can happen between two different
molecules or between different regions of the
same molecule
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Van der Waals InteractionsVan der Waals Interactions
• Due to the random movement of electrons
• Weak
• Short-lived
• Can occur in both polar and nonpolar molecules
• Only occur when molecules are very close
together
• Allows all molecules to be attracted to one
another
• Plays role in the shape of larger molecules
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Molecular ShapeMolecular Shape
• Every covalently bonded molecule has a
characteristic size and shape.
FOR IB BIO…the only thing about shape to
remember is:
• Biological Structure is related to function
– i.e. A molecule’s structure is directly related to
its “job”
• Molecules communicate via shape
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6CO6CO22 + 6H+ 6H22OO CC66HH1212OO66 + 60+ 6022
• Represented by chemical equations
– Reactants on the left
– Products on the right
– Some bonds are broken and reformed
– Mass is conserved in a reaction
• In a balanced chemical equation, the total # of
atoms of each element must be equal on both
sides of the equation
• Is this equation balanced?
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EquilibriumEquilibrium
• In some reactions, all of the reactants are
converted to products
• Most reactions, however, are reversible –
they can go in either direction
• CO2 + H2O H2CO3
• Eventually, equilibrium will be met.
– This is when the reaction is occurring in both
directions at the same rate
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Exergonic/Endergonic ReactionsExergonic/Endergonic Reactions
and Free Energyand Free Energy
• Free energy – energy that can be used to do work
• Exergonic reactions
– release free energy
– result in products with less stored energy than the reactants
– Reactants (high E) products (lower E) + E (free)
– C6H12O6 + 602 6CO2 + 6H2O + E
– Molecules are being broken down (catabolism)
• Endergonic reactions
– store free energy
– result in products with more stored energy then the reactants
– Reactants (lower E) + E (free) products (high E)
– 6CO2 + 6H2O + E C6H12O6 + 602
– Molecules are being built up (anabolism)
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Oxidation – Reduction ReactionsOxidation – Reduction Reactions
REDOXREDOX
• LEO the lion goes GER
• Loses e- oxidation, gains e- reduction
• Any time an ion is formed – redox reaction
• Example
– Na + Cl Na+
+ Cl-
• Na has lost e- and has been oxidized
• Cl has gained an e- and has been
reduced
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Redox in Covalent bondsRedox in Covalent bonds
• Redox rxns can also involve covalent bonding
• Atom can be reduced if it becomes bonded to a
highly eneg atom.
– i.e. it’s own e- are being pulled away from it
• Example
– C-H bond broken, H replaced with O, C-O
– Oxygen is highly e-neg
– Carbon has been oxidized
– Oxygen has been reduced
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Dalton’s Atomic TheoryDalton’s Atomic Theory
• We already have discussed this, but to
make it more clear the following 5 ideas
are the keys to make the Atomic Theory
more easy to identify
• 1. Elements are made of tiny particles
called atoms.
• 2. All atoms of a given element are
identical
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Dalton’s Atomic TheoryDalton’s Atomic Theory
• 3. The atoms of a given element are
different from those of any other element.
• 4. Atoms of one element can combine
with atoms of other elements to form
compounds. A given compound always
has the same relative numbers and types
of atoms.
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Dalton’s Atomic TheoryDalton’s Atomic Theory
• 5. Atoms are indivisible in chemical
processes. That is, atoms are not created
or destroyed in chemical reactions. A
chemical reaction simply changes the way
the atoms are grouped together.
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Describing SolutionsDescribing Solutions
• A solution is a uniform mixture
• Two types of parts
– Solvent –the dissolving agent
• Water is a great example (especially in cells)
– Solutes – are dissolved in the solvent
• Anything dissolved in a substance
• There can be many solutes in a given solvent
• Example – mix salt and water
– Water is the _____
– Salt is the _______
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Like Dissolves LikeLike Dissolves Like
• Polar vs. nonpolar
• Polar and nonpolar substances repel one
another
• So . . .
• Polar (and ionic) solutes will dissolve in
polar solvents
• Nonpolar solutes will dissolve in nonpolar
solvents
• Think oil (nonpolar) and vinegar (polar)
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Hydrophobic vs. HydrophilicHydrophobic vs. Hydrophilic
• HYDROPHOBIC
• Hydro = water
• Phobic = fearing
• Don’t dissolve in water
• Nonpolar substances
• OIL
• HYDROPHILIC
• Hydro = water
• Philic = loving
• Do dissolve in water
• Polar/ionic substances
• VINEGAR
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ReviewReview
• Like dissolves like
• Hydrophilic and hydrophobic, i.e. nonpolar and
polar molecules, literally repel one another
• All polar molecules are hydrophilic
• All ionic molecules are hydrophilic
• All nonpolar molecules are hydrophobic
• However
– Some molecules can be both hydrophobic and
hydrophilic (in different areas)
– Example – phospholipids
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Concentration of a SolutionConcentration of a Solution
• A measure of the amount of solute/solvent
• Lots of solute and/or low solvent = a high
concentration (represented by [x] )
• Aqueous solution – water is the solvent
– Very important to life
• Saturated solution – cannot dissolve any
more solute
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Dissociation into IonsDissociation into Ions
• To break into separate ions in solution
• Ionically bonded substances do this
– NaCl Na+
(aq) + Cl-
(aq)
• Covalently bonded substances don’t
dissociate into ions, with one exception
• Water is the “exception”
– H2O H+
+ OH-
• Note H+
and H3O+
are synonymous
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Acids and BasesAcids and Bases
• Acids
• H3O+
↔ H+
+ H2O
• H3O+
= Hydronium
• Acidity or alkalinity (bases) is actually a measure
of hydronium and hydroxide ions dissolved in a
solution
• BASES
• OH-
= hydroxide ion
• REMEMBER: NaOH ↔ Na+
+ OH-
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pH and lifepH and life
• Control of pH is very important to living things
(homeostasis)
• Example
– Human blood pH range generally 7.35 – 7.45
– Anything below 7 or above 7.8 can be deadly
• Buffers
– Weak acid/base that can neutralize small amounts of
another acid/base
– H2CO3 H+
+ HCO3
-
– Carbonic acid hydrogen ion + bicarbonate ion
