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Outline
 I. Why study Chemistry
 II. Elements
Chemistry
BioSci 105
Lecture 2
Reading: Chapter 2 (Pages 14-25)





Atoms
Isotopes
Periodic Table
Electrons
Bonding
 III. Bonds
 Covalent bonds
 Polarity
 Ionic bonds
 Hydrogen bonding
 IV. Water
 V. Acids and Bases
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Why study Chemistry?
Definitions and the Basics
 Chemistry is the basis for studying much of
biology
 Matter is anything that takes up spass and
have mass.
 The biology of the human body follows the
rules of physics and chemistry.
 Atoms are units of matter that cannot be
broken down into simpler substances.
 You need to understand enough about
chemistry to know what kind of things will
cross a membrane, and what are biological
compounds make up cells. What is a
protein?
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 An element is a “pure” form of matter
containing only one kind of atom.
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Elements in nature
Text page 22
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Atom
Table 2.1 Review of Subatomic Particles
 We used to think that an atom could not be
divided, now we know they are composed
of parts (particles):
 Protons – carry a positive charge
 Neutrons – have no charge (neutral)
 Electrons – carry a negative charge
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The Atom Cont.
 Protons and neutrons are in the center of the
atom
 Electrons orbit around the outer edge in
orbitals
 In each atom the # electrons = # protons
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Chemistry and Biology
Shell Model of Electrons
 Electrons can be visualized as residing in shells
around the nucleus.
 The first shell can have up to two electrons
 The second shell can have up to eight electrons
 The third, fourth … shells can have up to eight
electrons
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Figure 2.1c
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Periodic Table of Elements
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Figure 2.2
Particle Mass
 Proton = 1 amu
 Neutron = 1 amu
 Electron = negligible
The atomic number = the # of protons in an atom
Atoms have equal numbers of protons and electrons.
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Isotopes
Isotopes of Hydrogen
 Atoms with the same number of protons
but different numbers of neutrons are
called isotopes
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Radioactive Isotopes
 In 1896, Henri Becquerel placed a rock on
unexposed photographic plates inside a drawer.
The rock contained uranium.
 The isotopes of uranium emit energy.
 After a few days the plate had an image of the
rock.
 A co-worker, Marie Curie, named this
radioactivity. This is known as a radioisotope
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Radioisotopes in Medicine
Radioactive Isotopes
 Radioisotopes are isotopes that are
unstable, and become more stable by
emitting energy and particles
 In contrast, most isotopes are stable
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Isotopes in medicine
 PET scans (Positron-Emission Tomography)
 Patient is injected with a compound that is labeled
with an unstable isotope
 Cancer cells are growing faster and take up more
of the compound than normal cells
 Abnormal tissue takes up less of the compounds
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Figure 2.4
Atomic number
Atomic weight
The atomic weight = an average of the isotopes
For any element:
Mass number = round the atomic weight
Mass Number = (Number of Protons) + (Number
of Neutrons)
Number of Protons = Atomic Number
Number of Electrons = # Protons = Atomic Number
Number of Neutrons = Mass Number - Atomic Number
Number of Neutrons = Mass number - # of Protons
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Chemical Bonds
For Be:
Number of Protons
= Atomic Number = 4
Number of Electrons
= # of Protons = 4
Number of Neutrons
= Mass Number - Atomic Number
=9-4=5
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Electrons and Bonding
 Chemical bonds are unions between electron
structure from different atoms
 Molecules are when two or more atoms join
together. They can be the same element
(H2) or different elements (H2O)
 When different elements join the molecule is
referred to as a compound molecule
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Chemistry and Biology
 If the outer shell is full, then it is non-reactive
and stable = does not form chemical bonds.
 Incompletely filled outer orbital, then atom
reactive and will form chemical bonds.
 How many bonds it can form depends on
how many empty spots in outer shell
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Chemical Bonds
 Covalent bonds
 Ionic
 Hydrogen
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Figure 2.8
Covalent bonds
 Covalent bonds
 The strongest bonds
 They form when two or more atoms share
the electrons in their outer shells
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How many bonds can form?
 Each atom wants their outer shell filled.
 Hydrogen only has one electron in its shell
– wants two, so it can form one bond.
 Carbon has four electrons in outer shell,
wants eight, so it can form four bonds.
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Types of Covalent Bonds
Double Bond
Covalent bonds
•
•
Polar
Nonpolar
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1. Covalent Bonding
 When two atoms with unpaired electrons in
the outer most shell come together and
share electrons
 Each atom has an attractive force for the
other atoms unshared electrons, but not
enough to take it completely away
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Covalent Bonding
 Covalent bonds can be polar or nonpolar
 Nonpolar bonds the atoms have same pull on
the shared electrons (H2)
 Polar bonds – the atoms don’t equally share the
electrons (H2O)
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Polar Covalent Bond
Figure 2.11a
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Polarity
 Some atoms have a greater pull on shared
electron than other atoms
 The measure of this pull is electronegativity
 When a bond is made between atoms with
different electronegativities it is a polar bond
 The greater the pull the more electronegative
(remember that electrons are negative)
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Polarity Con’t
 Polar Covalent Bonding occurs with strong
electrophiles (electronegative): atoms with
nuclei that have a strong pull on electrons.
Common examples in biological molecules
include:
 Oxygen
 Nitrogen
 Sulfur
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Polarity
Water
Alcohol
H
 The oxygen side of water is slightly
negative and the hydrogen sides are
slightly positive
H
H
O
O
C
H
H
H
Aldehyde
Ketone
O
O
H2
C
H 3C
H2
C
C
C
H2
CH3
H3C
C
C
H2
H
Hydrocarbons
H2
C
H3C
H2
C
C
H2
H
H2
C
C
H2
CH3
C
H2
H
H
C
H
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H
H
H
S
N
CH3
CH3
Polar Groups
 Oxygen containing:
 Carboxyl = - COOH
 Hydroxyl (alcohol) = - OH
 Phosphates = -PO4
 Carbonyl
H
C
HC
CH
HC
CH
C
H
 Ketone = - CO
 Aldehyde = - CHO
 Nitrogen containing: Amino (-NH2)
 Sulfur containing: -SH
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Nonpolar compounds
Carboxyl
Alcohol
CH3CH2CH2OH
 Hydrocarbons – lots of carbons and
hydrogens bonded together
Ketone
Aldehyde
Ether
Hydrocarbons
CH3-O-CH2CH3
CH3CH2CH3
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Terminology
 Hydrophilic (water-loving) – polar molecules
that are attracted to water
 Hydrophobic (water-fearing) – nonpolar
molecules that are pushed aside by water
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Chemical formulas
 When we write compounds, we often write
them as a formula that tells how atoms
many of each element are present, but not
the way the molecule is put together.
 You often can determine the way the
molecule is put together by knowing how
many bonds each element can form.
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Examples
Examples
 How would you draw this compound?
 How would you draw this compound?
 H2O
 C4H10
H
O
H
H
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H
H
H
H
C
C
C
C
H
H
H
H
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Examples
Examples
 How would you draw this compound?
 C4H8
 How would you draw this compound?
 CO2
H
H
C
C
H
O
H
H
C
C
H
H
C
O
H
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Examples
2. Hydrogen Bonding
 How would you draw this compound?
 C2H4O
H
H
H
O
C
C
 Weak attraction between a hydrogen atom
with a partial positive charge and another
atom with a partial negative charge
(electronegative atom such as oxygen,
nitrogen, or sulfur).
H
H
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2. Hydrogen Bonding
2. Hydrogen bonds
 Individually weak, but many together can be
strong.
 Determines shapes of many biological
molecules including proteins and DNA
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3. Ionic Bonds
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Figure 2.11b
3. Ionic bond
 Ion = atom that has gained or lost electrons,
It no longer has a balance between protons
and electrons, it is positive or negative
charge
 Ionic bond is an association between ions of
opposite charge
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Figure 2.10
Chemical bonds
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Table 2.2
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Water – The Life Giving Molecule
Water’s Abundance
 Why are we so interested in finding evidence
of water on Mars?
 71% of Earth’s surface is water
 97.5% of the water is salt water
 Freshwater only accounts for 2.5% of water
 Only 0.53% is available to us to drink (rivers,
lakes, ground water)
 What would it mean if we did not find
evidence of water? Or if we find evidence?
Does it matter what form the water is?
 66% of the human body is water by weight
 75-85% of a cell’s weight is water
 Life exists here because water is abundant
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Water
 Water is polar and forms hydrogen bonds
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Hydrogen bonding
 Water is a great example of hydrogen bonding,
itis the hydrogen bonds that give water much of
its unique characteristics
H
O
H
O
H
H
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Exists in Three Forms
 Water exists in three forms
 Solid - Ice
 Liquid
 Vapor
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Ice
 As water approaches 0˚ C, freezing
temperature, the molecules slow down.
 Water forms more hydrogen-bonds at lower
temperature and forms a lattice structure
 The ice is less dense due to the lattice
structure and to the fact that there are less
molecules present than in the same volume of
liquid
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Ice caps
 The floating property of ice allows the Artic
Ocean’s ice cap to exist
 This is the habitat for polar bear and young
seals, as well as many other species.
 These ice caps are melting, as they melt the
habitat for these species shrinks.
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Properties of Water
 So what do all this mean?
 The polarity and ability to form hydrogen
bond give water its properties
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1. Water is an excellent polar solvent
 Because water is polar and forms hydrogen
bonds, it acts as a solvent for polar molecules
 Like dissolves in like, so polar molecules
dissolve in water
 There are four properties of water
 Water is considered the best polar solvent –
due in great part to its ability to form
hydrogen bonds with other molecules
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Why is this property important?
 Blood is approx 55% water so the fact that
water is a good solvent makes blood a good
way to transport things around.
 Cells are made up of mainly water, the water
keeps salts in your cells, blood and tissues in
solution.
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2. Water has Cohesion
2. Water has Cohesion
 Due to the hydrogen bonding, water has
cohesion (the water molecules cling together)
 What allows bugs to walk on water?
 Hydrogen-bonds create surface tension
 At the surface of water, where water meets air,
the water molecules are being pulled down with
a much greater force than they are being pulled
up towards the air
 Cohesion is the capacity to resist breaking
under tension
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Why is this property important?
 The cohesion of water allows blood move
easier in the blood vessels.
 Also is responsible for moving water in plants
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3. Water has High heat capacity
 It takes a great deal of energy to raise the
temperature.
 When you increase the temperature of
something, the molecule in it move faster,
hydrogen bonds keep the water molecules in
place so it takes lots of energy to break the
bonds and heat the water
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Why is this property important?
 Water in our bodies keep us at a constant
temp.
4. Water has High heat of Vaporization
 It takes a great deal of energy to make
water evaporate (change water from a
liquid to a gas).
 Hydrogen bonds must be broken in order
to change water from liquid to vapor
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Why is this property important?
 Sweat is mainly water, when we sweat the
body uses its heat to vaporize the water –
cooling us off.
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Acids and Bases
 We are already familiar with acids and bases
 Common acids:
 Lemon juice
 Sodas
 Vinegar
 Common bases:
 Ammonia
 Many household cleaners
 Bleach
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Acids and Bases
 Acids – Substances that donates hydrogen ions
when in solution
 HCl  H+ + Cl-
 Bases – Substances that accept hydrogen ions
when in solution
 NaOH  Na+ + OH-
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Ph Scale
 The strength of acids and bases is measured
using the pH scale.
 pH = -log10[H+]
 [H+] = conc in moles per liter
 It is inverse relationship:
 Higher the pH the lower the concentration of H+
 In solution:
 H+ + Cl- + Na+ + OH-  H2O + NaCl
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 Logarithmic:
 Each point increase in pH represents a ten-fold
decrease in H+ concentration.
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Ph Scale
Acids and Bases
 Scale from 0 – 14
 0 is the most acidic
 14 is the most basic
 7 is neutral (pure water)
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Table 2.3
Biological Fluids
 Blood – pH 7.35 Changes in pH of ± 0.1
can damage cells, pH of 7.8 can be lethal
 Biological fluids have buffers to keep the
pH stable.
 Most biological fluids are between 6 – 8
 Stomach fluid – pH of under 2
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Buffers
 Buffers resist pH changes because they
are chemicals that can take up excess H+
or OH Our body wants to keep its fluids at an
even pH.
Buffers
 For example when CO2 enters the blood it
combines with H2O to form carbonic acid
(H2CO3).
 This weak acid dissociates to form H+ and
bicarbonate ion (HCO3-)
 Blood contains buffers that are weak acids
that can dissociate into ions.
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Acids in the Environment
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Acid Rain
 Read the section on page 25
 The two main sources of acid rain are H2SO4
and HNO3
 Most of H2SO4 pollution comes from electric
power plants
 Most of HNO3 pollution comes from cars,
buses etc
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Effects of Acid Rain
 Acid rain acidifies lakes and streams
 More acidic water leaches more heavy
metals from the soil than normal water.
 Declining frog and fish populations may be
due more to increased metal
concentrations in the water than due to the
acidic water itself
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Important Concepts
 Reading for next lecture: Chapter 2 (pages 25 –
36)
 What are the three particles of an atom, where are
they located, what is their charge, and mass.
 Be able to determine how many bonds each
element can form.
 Be able to recognize if a molecule is drawn
correctly.
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Important Concepts
 Be able to read the periodic table to determine the
number of protons, neutrons and electrons in the
atoms of all the biologically important elements.
 What are the three most common elements in the
human body
 Be able to draw the atom of any biologically
important element, with the correct number of
protons, neutrons, and electrons. Be able to draw
the electrons in their correct shell.
Important Concepts
 Be able to describe the types of chemical bonds
 Be able to draw a water molecule and hydrogen
bonding between water molecules
 Be able to describe the four properties of water
and their importance in living organisms.
 Understand the pH scale
 Be able to identify polar and nonpolar molecules
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Important Concepts
 Understand the causes and effects of acid rain,
know the chemicals that cause acid, the effects
on the environment, and the human health
effects associated with acid rain (see page 25)
 What are three electronegative elements found
in biological molecules
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Definitions
 Matter, Atom, Element, Isotopes, Radioisotopes,
Chemical bonds, single bond, double bond,
Molecules, Compound Molecules, Ion, Ionic bond,
Covalent bond, Nonpolar bonds, Polar bonds,
electronegativity, Hydrogen bond, Hydrophilic,
Hydrophobic, Cohesion, acid, base, buffers,
logarithmic, inverse, pH, solvent, solute, solution
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