Transcript Chemistry Pre-Quiz!

Chemistry
Pre-Quiz!
Question #1
Biochemistry is the
study of
____________.
Question #2
An atom consists of
protons, neutrons, and
____________.
Question #3
The molecular formula
for water is
____________.
Question #4
Which is the smallest?
 Organism
 Cell
 Atom
 Compound
Question #5
Glucose is an example
of a monosaccharide.
True or False?
Chapter 2
The
Chemistry of
Life
Why do biologists need
to study chemistry?
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All life forms are MATTER and all matter is
made up of chemical parts.
Matter = anything that has a MASS and
VOLUME that can be measured.
Examples of Matter (living and nonliving)
All matter has both physical and chemical
properties.
Physical Properties
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Can be observed and measured without
permanently changing the identity of the
matter
Observed by the senses - taste, smell,
shape, color, texture, melting and boiling
points.
Chemical Properties
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Chemical properties can change a substance
into a new substance through a chemical
reaction.
New substance is created and permanently
altered!
Rusting (oxidation) and combustion
(flammability)
Examples of chemical reactions
Levels of Organization
(Largest
Smallest)
Organism
Organ Systems
Organs
Tissues
Cells
Cell Organelles
Molecule/Compounds
Atoms
Ch. 2- Section 1
ATOMS
Smallest, indivisible unit of matter.
 Extremely small!
(100 MILLION atoms lined up end to end
would only equal 1 cm in length!)
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*However, an atom is made up of smaller
parts called SUBATOMIC PARTICLES.
ATOMS (continued)
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Center of Atom = Nucleus
Subatomic Particles = Protons, Neutrons,
Electrons
Protons - positively charged particle (+1)
Neutrons- particle with NO CHARGE; neutral
(0)
Electrons- Negatively charged particle (-1)
ATOMS (continued)
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Nucleus contains protons AND neutrons.
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Electrons circle around the nucleus in
electron clouds or orbitals or valence
shells.
ATOMS (continued)
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1)
2)
3)
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Every atom has an ATOMIC NUMBER. The
atomic number tells you a lot of information.
This number tells you the type of atom you have.
(Element name)
This number ALWAYS tells you the # of Protons
in the atom.
The atomic # is usually the number of neutrons
and electrons, also. (not always!)
Example: Atomic Number of 6 =
6 Protons, 6 Neutrons in the nucleus of the
atom.
Use Periodic Table of Elements to look up
element name for atomic number 6 = Carbon
Elements
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Every element consists of only 1 type of
atom.
Every element has a DIFFERENT
atomic number and its own symbol on
the Periodic Table (PT).
109 elements on the PT, 90 are found in
nature, 19 were formed in a laboratory by
scientists.
Using the Periodic Table
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Below the Symbol (C) you’ll find the
element’s atomic mass.
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Atomic Mass = Sum of protons +
neutrons
Drawing Atomic Structures
1)
2)
3)
4)
5)
6)
Look up element on the Periodic Table.
Write down Atomic Number.
Draw Nucleus with 3 orbitals surrounding it.
Inside nucleus, write # of protons (p+).
Inside nucleus, write # of neutrons (n0).
Fill in orbitals with correct number of electrons.
*First orbital gets only 2 electrons.
*Additional electrons, fill 2nd orbital.
(up to 8 e-)
*More electrons will go into 3rd orbital.
(up to 8 e-).
Drawing Atomic Structures
Example: CARBON
6 p+
6 n0
PRACTICE!
Drawing Atomic Structures
Draw the following elemental atoms.
•
Oxygen
•
Nitrogen
Chemical Elements
& Compounds
Remember!
The atomic number of an element NEVER changes!
IT IS ALWAYS THE # OF PROTONS IN THE
NUCLEUS!
However, the number of NEUTRONS can vary from
one atom of that element to the next.
Isotopes
Examples:
13C 14C
All are Carbon atoms, just with different atomic
MASSES! (same atomic #)
Atoms of the same element, same number of
protons, but different # of neutrons are called
isotopes.
Radioactive Isotopes
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The nuclei of some atoms are unstable and will
from time to time break down, releasing matter
and energy that we call radiation.
Atoms that emit (give off) radiation are said to
be RADIOACTIVE.
Radioactive Isotopes have many uses for
living organisms, such as diagnosing and
treating diseases, sterilize food, and measuring
the ages of things, like trees or rocks.
Radioactive Isotopes (continued)
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Examples:
Iodine - 131 (131I): injected into humans to study
the function of the thyroid gland. Can be seen
through special equipment that picks up on the
radiation energy given off by this isotope as it
travels through the body.
Carbon-14 (14C): used to treat brain tumors and
track the ages of trees and fossils.
Chemical Bonding
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When individual atoms interact, they form
chemical bonds.
When chemical bonds form between 2 or more
atoms, your result is a chemical compound.
However, atoms combine in a certain way and
follow a specific set of rules!
The rules affect the electrons in the outermost
orbital, or VALENCE ELECTRONS.
VALENCE ELECTRONS
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When the outermost orbital is full with the
maximum number of electrons, the atom is very
stable (happy!). These atoms are unreactive and
do not combine with other atoms to form
compounds.
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When the outermost orbital is NOT full, it will try
and fill its orbital by combining with other atoms.
To become stable, an atom will either GAIN,
LOSE, or SHARE electrons. An atom will bond
with another atom if the bond gives both atoms
complete valence energy levels.
Two major types of bonds
that can form compounds:
1.
2.
IONIC BONDS
COVALENT BONDS
IONIC BONDS
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DEFINITION: A bond that completely
transfers electrons from one atom to
another.
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Electrons are completely given up or
taken in order for atoms to have a
completely full outermost orbital.
Remember- that’s when atoms are
most happy and stable!
IONIC BONDS (cont.)
11 p+
11 n0
Na
17 p+
17 n0
Cl
IONIC BONDS (cont.)
Happy!
NaCl
11 p+
11 n0
Na
11 p+
+ 10 e+1
1+
17 p+
17 n0
Cl-1
17 p+
+ 18 e1-
Happy!
Ionic Bonds (cont.)
When electrons are gained or lost, the elements take
on a charge. The charge will be either positive or
negative based on gaining or losing electrons.
These charged elements are called IONS.
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http://www.youtube.com/watch?v=Ftw7a5ccubs&featur
e=related
COVALENT BONDS
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DEFINITION: A bond that shares
electrons from one atom to another.
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Electrons are NOT completely given up
or taken, BUT ARE SHARED. This
allows for atoms to have a completely
full outermost orbital.
COVALENT BONDS
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Example: H2O
Ionic & covalent bonding video
More covalent
bonds!
Chapter 2 – Properties of Water
Properties of Water
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Polar covalent bonds- e- are shared,
however, electrons are pulled more towards
the oxygen, due to the more protons
(+ charges) in the nucleus.
Water is POLAR due to this PULL, or uneven
distribution of electrons.
Hydrogen bonds (weak bonds)
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Cohesion- attraction between molecules of
the SAME substance
(Ex. Water beads up due to the attraction
between similar water molecules- “Surface
tension” on penny)
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Adhesion- attraction between molecules of
the DIFFERENT substances
(Ex. Adhesives like glue, tape, etc.)
Acids/ Bases/ pH Scale
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Water can break apart into ions
H2O
H+
Water
Hydrogen
Hydroxide
Ion
Ion
+
OH-
The pH scale measures the
concentration of H+ ions in a solution.
Acids/Bases
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pH scale ranges from 0-14.
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Any reading BELOW 7 on the pH scale is an
ACID. (HIGH concentration of H+ ions)
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Any reading ABOVE 7 on the pH scale is an
BASE. (LOW concentration of H+ ions)
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Acids (Acidic solutions) – lower the pH, the
STRONGER the acid.
Bases (ALKALINE solutions) – higher the pH,
the STRONGER the base.
Buffers
Any pH level near 7, is consider
NEUTRAL.
 Buffers can be added to strong
acids/bases to prevent strong changes
in the pH levels.
* Very important for maintaining
blood pH in humans
(HOMEOSTASIS)
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