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biochem notes

 The study of chemistry begins with
the basic unit of matter, the atom.
 The Greek philosopher Democritus
called the smallest fragment of matter
the atom, from the Greek word
Atoms (cont.)
 Placed side by side, 100 million atoms
would make a row only about 1
centimeter long.
 Atoms contain subatomic particles that
are even smaller.
Atoms (cont.)
 What three subatomic particles make
up atoms?
Atoms (cont.)
 The subatomic particles that make up
atoms are
 protons
 neutrons
 electrons
Atoms (cont.)
 Smallest property of an element that still
has the properties of that element
 “The building blocks of matter”
 Atoms are made of smaller (subatomic)
particles arranged in a particular way
 p+ (proton)
 n° (neutron)
 e- (electron)
Atoms (cont.)
 Protons and neutrons have about the
same mass.
 Protons are positively charged particles
 Neutrons carry no charge (◦).
 Strong forces bind protons and
neutrons together to form the nucleus,
which is at the center of the atom.
Atoms (cont.)
 The electron is a negatively charged
particle (−) with 1/1840 the mass of a
 Electrons are in constant motion in the
space surrounding the nucleus (ecloud).
Atoms (cont.)
 The subatomic
particles in a
helium atom.
Atoms (cont.)
• Electrons are attracted to the
positively charged nucleus but remain
outside the nucleus because of the
energy of their motion.
• Because atoms have equal numbers of
electrons and protons, and because
these subatomic particles have equal
but opposite charges, atoms are
Atoms (cont.)
Atomic number- # of p+ AND electrons in an atom
Mass number- total # of p+ + n° in an atom
# of neutrons- mass number- atomic number
Ion- charged atom (can gain or lose electrons)
 If an atom gains e-, does it have a positive or
negative charge?
 What if the opposite is true?
 LEP-GEN (Lose Electrons Positive, Gain Electrons
 Where can the characteristics of all atoms be
 Atomic #?
 Atomic mass?
 # of neutrons?
More Practice
 Atomic #?
 Atomic mass?
 # of neutrons?
 Substances that can’t be broken down into simpler
 Pure substance made of only ONE type of ATOM
 Represented by one or two letter symbol (Ex. C, H,
O, Al, Fe, He, Ga, Pt, Au)
 Noble gases
 Elements that can exist alone (don’t combine with other
 Ex. He, Ne, Ar
 Mike Stanfill, Private Hand - Flash Animation - The
Elements, by Tom Lehrer
Elements (cont.)
 Sodium
 Reaction of Sodium with Water
 Mercury
Elements (cont.)
• More than 100 elements are known, but
only about two dozen are commonly
found in living organisms.
 Atoms of the same element that have
different #’s of n°
 Ex. C-14 and C-12
 Both have 6 p+
 C-14 (8 n°), C-12 (6 n°)
Isotopes (cont.)
 Isotopes are identified by their mass
 For example, carbon has three
isotopes—carbon-12, carbon-13, and
carbon-14. Each isotope has a
different number of neutrons.
Isotopes (cont.)
 Q- How are all of the isotopes of an
element similar?
Isotopes (cont.)
 A- Because they have the same number
of protons and electrons, all isotopes
of an element have the same chemical
Isotopes (cont.)
Isotopes of Carbon
6 electrons
6 protons
8 neutrons
Isotopes (cont.)
 Radioactive Isotopes
 Some isotopes are radioactive, meaning
that their nuclei are unstable and break
down at a constant rate over time.
 Although the radiation these isotopes give
off can be dangerous, they have important
scientific and practical uses.
Isotopes (cont.)
• Radioactive isotopes can be used:
 to determine the ages of rocks and
 to treat cancer.
 to kill bacteria that cause food to spoil.
 as labels or “tracers” to follow the
movement of substances within an
Chemical Compounds
 Chemical Compounds
 In nature, most elements are found
combined with other elements in
 A chemical compound is a substance
formed by the chemical combination of
two or more elements in definite
 The physical and chemical properties of a
compound are different from the
elements from which it is formed.
Chemical Compounds
• Scientists show the composition of
compounds by a kind of shorthand
known as a molecular formula.
• Water, H2O, contains two atoms of
hydrogen for each atom of oxygen.
• The formula for table salt, NaCl,
indicates that sodium and chlorine
combine in a 1 : 1 ratio.
Molecular Formulas
 Shows numbers of molecules and atoms
 Molecules = coefficient (the number in front) of
molecular formula (ex. 3H2O = 3)
 Atoms = # of atoms in compound
(ex. H2O = 2 Hydrogens, 1 Oxygen = total of 3
atoms in this compound)
 Total # of atoms = (Coefficient) x (# of atoms)
 How many molecules and atoms do each of the
following have?
 3 H2O
 2 C6H12O6
 4 NaCl
 CO
Structural formulas
 Shows how the atoms are connected in
a compound. Each element requires 1-4
lines to connect it to other atoms (see
periodic table)
 Ex. C2OH6
 Is there a problem in drawing this?
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 Have same molecular formula but
different structural formulas
 C3H8O
Energy shells
 Located in the e- cloud
 1st shell holds a maximum of 2 e All other shells hold up to 8 e-
Electron Configuration
 Drawing of how many e- are in each orbital.
 Must know atomic number so know how many
protons (and therefore how many electrons)
an element has
 Ex. O has atomic # of 8 therefore has 8 e Electron configuration = 2) 6)
 Ex. Al has atomic number of 13
 Electron configuration = 2) 8) 3)
 Shows how elements bond together by
sharing e Several steps involved in determining
 What is an e- orbital?
 How many e- can fit in each orbital?
 What do we have to know before we can
determine valence?
Valence electrons
 Number of e- in the outermost energy
 Ex. O2 has _______ valence e-
Valence (cont.)
 Where are the valence e- located?
 Answer: In the LAST orbital!!!
 _______ +7 2)
 _______ +5 2)
 Each orbital wants to be FULL to be happy!!
 To determine valence, always take the path
of least resistance (always choose the
lowest number to gain or lose to make the
orbital full)
Valence (cont.)
 If an atom loses e-, does it become
more positive or more negative?
 This positive or negative will always be
the sign in front of the valence
 The valence number is found by
figuring out how many e- must be lost
or gained to make the atom happy
(have the last orbital full!)
Valence (cont.)
In summary, to find valence:
1. Look up the atomic number (# of p+)
2. Determine the electron configuration
3. Determine the # of valence e- (# of electrons
in last shell)
4. Determine if it is + or – valence (if it gains e- it
becomes more negative and vice versa)
5. Determine valence by how many e- it would have
to lose or gain to be happy
Valence (cont.)
 Want an easier way to find valence?
 Label your periodic table!
Chemical bonds
 Forces that hold atoms together
 Can be single, double, or triple bond
 Depends on how many pairs of electrons
are shared between elements
 Single bond
 Double bond
 Triple bond
Bond Lines
 All elements have the ability to bond
with other element(s)
 Can have anywhere between 1 and 4
bond lines
 How do we know how many bond lines
each element has?
 Answer: the number of bond lines is the
same of the valence of the element
(without the + or – sign)
Bonding (cont.)
 Ex. Cl has a valence of -1, so it has one
bond line.
 Ex. C has a valence of ±4, so it has four
bond lines.
 Ex. N has a valence of -3, so it has
three bond lines.
Bonding rules
 Each element must have the correct #
of bond lines attached to it
 # of bond lines is determined by
 Ex. O2 has valence of -2, so it has
_________ bond lines.
 Ex. C has a valence of ± 4, so it has
_________ bond lines.
Bonding rules (cont.)
 All coefficients and subscripts must be
 Ex. 2C2H6- Must draw two molecules
each having 2 carbons and 6 hydrogens
Changing molecular formulas
to structural formulas
 All bond lines, subscripts, and
coefficients must be satisfied.
 Ex. H2O
Changing structural formulas
to molecular formulas
 All coefficients and subscripts must be
 Ex. H – O – H ―> H2O
 Ex. H – Cl H – Cl ―> 2HCl
Types of bonds
 Covalent
 Chemical bonding where pairs of electrons are
 Strongest type of bond (gets stronger as more
pairs of electrons are shared)
 Polar covalent bond
 Present in water molecules
 Because electronegativity difference between O
and H, a bond forms where the O has a partial
negative charge and H a partial positive charge
Types of bonds
 Ionic bond
 Weak bond formed between two
oppositely charged ions
 Type of bonds in salt (NaCl)
 Hydrogen bond
 Caused when hydrogen and an
electronegative atom bond
 Weakest type of bond
Organic compound
 Must contain carbon (C)
 Usually associated with living things
 Ex. Carbohydrates, lipids, proteins,
nucleic acids
 What is inorganic????
Chemical Bonds
 What are the two main types of
chemical bonds?
Chemical Bonds
 The main types of chemical bonds are:
 ionic bonds
 covalent bonds
Chemical Bonds
 Chemical Bonds
 The atoms in compounds are held together
by chemical bonds.
 Bond formation involves the electrons
that surround each atomic nucleus.
 The electrons that are available to form
bonds are called valence electrons.
Chemical Bonds
 Ionic Bonds
 An ionic bond is formed when one or more
electrons are transferred from one atom
to another.
 An atom that loses electrons has a
positive charge.
 An atom that gains electrons has a
negative charge.
 These positively and negatively charged
atoms are known as ions.
Sodium atom
Chlorine atom (Cl)
Chloride ion (Cl-)
Sodium ion (Na+)
- 11
- 17
Chemical Bonds
 Covalent Bonds
 Sometimes electrons are shared by atoms
instead of being transferred.
 Sharing electrons means that the moving
electrons actually travel in the orbitals of
both atoms.
Chemical Bonds
 A covalent bond forms when electrons
are shared between atoms.
 When the atoms share two electrons, the
bond is called a single covalent bond.
 When atoms share four electrons it is
called a double bond.
 When atoms share six electrons it is
called a triple bond.
Chemical Bonds
 The structure that results when atoms
are joined together by covalent bonds
is called a molecule.
 A molecule is the smallest unit of most
Chemical Bonds
 In a water
molecule, each
hydrogen atom
forms a single
covalent bond
with the oxygen
Chemical Bonds
 Van der Waals Forces
 When molecules are close together, a
slight attraction can develop between the
oppositely charged regions of nearby
 Chemists call such intermolecular forces
of attraction van der Waals forces, after
the scientist who discovered them.
Chemical Bonds
• Although van der Waals forces are not
as strong as ionic bonds or covalent
bonds, they can hold molecules
together, especially when the
molecules are large.
Chemical Bonds
 For example, van der
Waals forces form
between the molecules on
the surface of a gecko’s
foot and the molecules on
the surface of the wall.
 The combined strength of
all the van der Waals
forces allows the gecko to
grip the wall.
 The particles that move around the
nucleus of an atom are called
 neutrons.
 protons.
 electrons.
 isotopes.
 The atomic number of a carbon atom is 6.
How many neutrons does the isotope
carbon-14 have?
 12
 14
 Which of the following statements about
the three isotopes of carbon is true?
 They are all radioactive.
 They have different numbers of electrons.
 They have the same chemical properties but
differ in atomic mass.
 They have the same number of protons and
 A chemical compound consists of
 electrons mixed with neutrons.
 two or more elements combined in a definite
 two or more elements combined in any
 at least three elements combined by ionic or
covalent bonds.
 Van der Waals forces are the result of
 unequal sharing of electrons.
 ionic bonds.
 the bonding of different isotopes.
 the chemical combination of sodium and