Transcript Chemical bonds - Chipola College

Chemical bonds
An introduction to chemistry
Compounds and chemical
change
• Atom - smallest elemental
unit
• Molecule
– smallest particle still
retaining the characteristic
chemical properties of a
substance
– Examples:
• oxygen, hydrogen gas diatomic molecules
• Ozone - triatomic oxygen
molecule
• Noble gases: helium, neon
- “monatomic” molecules
Chemical reactions
• Formation and/or
breaking of chemical
bonds to form new
molecules (products)
from old ones
(reactants)
• Chemical energy internal bonding
potential energy
• Chemical equation symbolic summary of
chemical reaction
• Magnesium is an alkaline
earth metal that burns
brightly in air, releasing
heat \and light. As
chemical energy is
released, a new chemical
substance is formed. The
new chemical material is
magnesium oxide, a soft
powdery material that
forms an alkaline solution
in water (called milk of
magnesia).
• A chemical equation is a shorthand
representation of a chemical reaction.
– The substances that are changed in to reaction are
called the reactants.
– The substances that are formed in the reaction are
called the products.
– chemical reactions and energy flow can be
explained by the making or breaking of chemical
bonds.
-- chemical bonds can be explained in terms of
changes in the electron structure of atoms
• Valence electrons are those electrons that
occur on the outer shell of atoms.
– These are the electrons that are responsible for
chemical reactions
– Inner electrons are in lower energy levels and the
orbitals are filled.
• These electrons are therefore not available for
interactions with other electrons.
• The Octet Rule
– Atoms attempt to acquire an outer orbital with eight
electrons through chemical reactions (octet rule)
– This gives them an outer shell configuration like
their nearest noble gas and therefore they become
stable.
– From the family number of the representative
elements, you can determine the number of valence
electrons, and therefore the number of electrons
necessary to gain the stable configuration
– Sodium is in Group IA and therefore has 1 electron
in its outer shell.
• If sodium loses one electron it becomes Na+ and
has 8 electrons in its outer shell.
• It then has the electron configuration of Ne
(neon) and has a filled outer shell configuration.
Valence electrons and ions
• Outer electrons
determining the
chemical properties of
an atom
• Octet rule
– Atoms attempt to acquire
an outer shell of eight
electrons
– Electrons can be
gained/lost/shared in the
process
• Example: sodium (Na)
Chemical bonds
• Attractive forces holding
atoms together in
compounds
• Can be described in terms of
molecular (delocalized) or
atomic (localized) orbitals
• Covalent
Three types:
• Ionic
• Metallic bonds
– Electrons transferred
between atoms
– Electrostatic force =
binding force
– Octets achieved through
sharing electrons
– Typically between
nonmetallic elements,
r.h.s of periodic table
– Outer electrons move
freely throughout metal
– “Electron gas” within rigid
lattice of metal atoms
– Conduct heat and
electricity well
Ionic bonds
• Chemical bond of
electrostatic attraction
• Form crystalline solids
with orderly geometric
structure
• Example: NaCl
• Na loses; Cl gains
• No single NaCl
molecule, per se
• Sodium chloride crystals are composed of
sodium and chlorine ions held together by
electrostatic attraction. Each sodium ion is
surrounded by six chlorine ions, and each
chlorine ion is surrounded by six sodium ions. A
crystal builds up like this, giving the sodium
chloride crystal a cubic structure.
• You can clearly see the cubic structure of these
ordinary table salt crystals because they have been
magnified about ten times.
Energy and electrons in ionic
bonding
• Reaction energy
released = heat of
formation
• Divided conceptually
into half-reactions
Electron transfer rules
• Electrons lost/gained to
form closed octets
• Number gained =
number lost
• Two rules for keeping track of electrons in ionic
bonding reactions.
–Ions are formed when atoms gain or lose
electrons to achieve a noble gas configuration
–The number of electrons that are lost must
equal the number of electrons that are gained.
– Electrons are not created or destroyed in a
chemical reaction
Ionic compounds and
formulas
Formulas
•
•
List elements in compound and
their proportions
Proportions decided by electron
gain/loss
Ionic compounds
•
•
•
•
Characterized by ionic bonds
White, crystalline solids soluble
in water
Families IA and IIA lose
electrons and form positive ions
Families VIA and VIIA gain
electrons to form negative ions
Covalent bonds
• Chemical bonds formed by sharing
pairs of electrons
• Electrons shared to form octets, ideally
• Overlap of shared electron clouds
between nuclei yields net attraction
• Atoms within covalent compounds are
electrically neutral, or nearly so
• Covalent Bonds
– A covalent bond is a chemical bond that is formed
when two atoms share a pair of electrons.
– H. + H.  H:H
– Covalent Compounds and Formulas
• Since a pair of electrons is shared in a covalent bond, the
electrons move throughout the entire molecular orbital.
• In the above example, since both hydrogen share the
electron pair, each hydrogen has a filled valence shell,
since it has the electron configuration of helium.
• Compounds that are held together by covalent bonds are
called covalent compounds.
• Covalent compounds form from atoms on the right side of
the periodic table
• Families IVA through VIIA, the number of unpaired
electrons (number of covalent bonds) is eight minus
Covalent compounds and
formulas
• Covalent compound held together by
covalent bonds
• Electrons shared in
covalent bonds
• Electron dot
representation
– Bonding pairs shared
– Lone (non-bonding) pairs
not shared
Multiple bonds
• Sharing of more
than one electron
pair
• Examples
– Ethylene - double
bond
– Acetylene - triple
bond
• Acetylene is a hydrocarbon consisting of two carbon atoms and
two hydrogen atoms held together by a triple covalent bond
between the two carbon atoms. When mixed with oxygen gas
(the tank to the right), the resulting flame is hot enough to cut
through most metals.
– Coordinate Covalent Bonds
• A coordinate covalent bond is one in which the
electron pair comes from one atom.
H+ + NH3  NH4+
H
H
..
..
H+
+
:N:H  H:N:H
..
..
H
H
• In this example both electrons from the new
covalent bond come from the lone pair (nonbonding electrons) around the nitrogen.
Bond polarity
• Result of unequal
sharing of electrons
• Electronegativity
– Measure of an atom’s
ability to attract electrons
– Differences:
• 1.7 or greater - ionic
• 0.5-1.7 - polar covalent
• Less than 0.5 - covalent
Electronegativities
Composition of compounds
• Millions of different combinations of over 90
elements
• Common names
– Often related to historical usage (baking soda,
washing soda,…)
– Difficult to relate to actual molecular composition
• Modern approach - systematic sets of rules
– Different for ionic and covalent compounds
– One common rule - “-ide” means compound
contains only two different elements
• These substances are made up of sodium and some
form of a carbonate ion. All have common names with
the term "soda" for this reason. Soda water (or "soda
pop") was first made by reacting soda (sodium
carbonate) with an acid, so it was called "soda water."
Ionic compound names
• Name of metal (positive) ion
first; then nonmetal
(negative) ion
• Many elements have
variable charges
• Historical suffix usage
– “-ic” for higher of two;
– “-ous” for lower
• Modern approach
– English name of metal
followed by Roman
numeral indicating
charge
• Ionic Compound Names
– Ionic compounds that are formed from metal ions
are named by naming the metal ion (electropositive
ion) first, followed by the nonmetal (electronegative
ion)
– The ending of the nonmetal is changed to end in ide
– When a metal can have various oxidation states the
oxidation state is give by roman numerals in
parenthesis after the name of the metal. Older
naming added –ic & -ous to the end of the name of
the metal, for the higher and lower of two possible
charges.
-- some ionic compounds have polyatomic ions,
thereby containing 3 or more elements. ( Metal
Ionic compound formulas
• Two rules
– Write symbol for
positive ion first
followed by negative
ion symbol
– Assign subscripts to
assure compound is
electrically neutral
• Example: Calcium
chloride
• A battery and bulb will tell you if a solution
contains ions.
Covalent compound names
• Molecular composed of two or
more nonmetals
• Same elements can
combine to form a
number of different
compounds
Two rules
• First element in formula
named first with number
indicated by Greek
prefix
• Stem name of second
element next; Greek
prefix for number;
ending in “-ide” for two
elements)
• Covalent Compound Names
– Names for covalent compounds uses Greek prefixes to
indicate numbers of atoms of each element
• The first element in the formula is named first with a prefix
indicating the number of atoms if the number is greater
than one.
• The stem name of the second element in the formula is
named next, with a prefix used with the stem name if two
elements can form more than one compound
– The suffix –ide is again used.( either ionic or covalent)
• Example
– CO = carbon monoxide
– CO2 = carbon dioxide
• Covalent Compound Formulas
– The systematic name tells you the formula
– Formulas indicate how many atoms of one element
combine with atoms of another element.
– The number of covalent bonds that an atom can
form is called its valence.
– Lone pairs create the possibility of creating
coordinate covalent compounds.
• Once you understand chemical names and formulas, you can
figure out what chemical compounds are contained in different
household products. For example, (A) washing soda is sodium
carbonate (NA2CO3) and (B) oven cleaner is sodium hydroxide
(NaOH), which is also known as lye.
Covalent compound formulas
• Examples: carbon dioxide,
carbon tetrachloride
• Valence
– Number of covalent
bonds an atom can form
– Hydrogen valence = 1
– Oxygen = 2; single and
double bonds
– Nitrogen = 3; single,
double and triple bonds
– Carbon = 4 - single,
double and triple bonds
• As you can see by
studying these two
charts, there is a
relationship between the
number of bonding
electron pairs and the
number of lone electron
pairs and the shape of a
molecule.