Download Biology\Ch 2 Chemistry

Unit 2: Biology (Chemistry Concepts)
Objectives: I can …
•Identify common elements or compounds by their symbols or formulas.
•Write and read chemical equations and understand how/why reactions occur
•Differentiate between atomic mass & atomic number; between atoms,
molecules, ions, & isotopes; between buffers, acids & bases; between
protons, electrons, & neutrons; between carbohydrates, lipids, proteins,
nucleic acids, & enzymes; between solute, solvent, solution, suspension,
concentration, etc.
•Distinguish between types of bonds (Ex: covalent, ionic, hydrogen, etc) and
their characteristics.
•Relate the unique properties of water to polarity, hydrogen bonds, gravity,
and molecular mass
•Analyze the relationship between states of matter and kinetic energy.
•Compare & contrast the types of energy
Vocabulary:
Matter, mass, atom, nucleus, neutron, proton, electron,
element, atomic number, atomic mass, chemical symbol,
isotopes, compounds, chemical formula, chemical equation,
ions, ionic bonds, covalent bonds, molecules, structural
formula, polarity, cohesion, adhesion, mixture, solution,
isotonic, hypotonic, hypertonic, solute, solvent, concentration,
suspension, colloid, hydronium ion, acid, hydroxide ion, base,
alkaline, pH, energy (potential, kinetic), reactants, products,
activation energy, exothermic, endothermic, monomer,
polymer, carbohydrate, lipid, protein, enzyme, nucleic acid,
substrate
matter - Occupies space and has mass
mass - The amount of matter, or material, in an object (How
is this related to weight?)
atom - Basic unit of matter. Each type of atom has a
unique number of subatomic particles (called protons,
neutrons, and electrons) that determine its structure and
behavior.
Subatomic particles:
1) neutrons - have no charge, weigh approx. 1 amu, are in
the nucleus (center) of the atom
2) protons - have a +1 (positive) charge, weigh 1 amu, and
are also in the nucleus
3) electrons - have a -1 (negative) charge, weigh almost
nothing, and move extremely rapidly at great distances
from the nucleus.
There is an equal number of protons and electrons in an
atom so an atom has no overall charge.
Elements are a single type of atom. An element may be
more than one atom of the same type. (Ex: elemental
oxygen is naturally found as O2 - two atoms of oxygen)
Each element is listed in the periodic table and is given a
chemical symbol, an atomic number, and an atomic
mass number.
• Chemical symbol - a unique abbreviation (similar to
initials) for each element. This is used in formulas and
equations.
• Atomic number - the number of protons (positive charge)
in an atom. Since atoms always have the same number
of protons and electrons, this also tells us how many
electrons are present.
• Atomic mass - the number of protons plus the number of
neutrons (This is not usually a whole number since
neutrons weigh slightly more than protons.)
An Element in the Periodic Table
Section 2-1
6
C
Carbon
12.011
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Section:
Examples of commonly used chemical symbols:
C = carbon
Ca = calcium
Na = sodium (natrium)
K = potassium (kalium)
H = hydrogen
He = helium
Cl = chlorine
Mg = magnesium
*Note: Only the first letter is capitalized for a chemical symbol. If we have
2 capitalized letters next to each other, we have 2 separate atoms and
are looking at a chemical formula.
Example:
Co = cobalt (a single element)
CO = carbon monoxide (a molecule with 2 types of atoms)
To determine the number of atoms of each element in a chemical
formula, we look at the subscripts. If no subscript, there is one atom of
that element.
Ex: CO = 1 carbon, 1 oxygen
CO2 = 1 carbon and 2 oxygen atoms
Chemical formulas show how many and what kinds of
atoms are present in a molecule or compound. Ex:
MgCl2 shows there are 2 chlorine atoms for every
magnesium atom.
C6H12O6 shows there are 6 carbon atoms, 12 hydrogen
atoms, and 6 oxygen atoms.
CO shows there is 1 carbon and 1 oxygen (Co is cobalt)
Molecules are atoms held together because they share
electrons. This is a covalent bond.
Ionic compounds are held together because 1 atom has a
stronger pull for electrons than the other so it pulls the
electron(s) over making it slightly negatively charged and
the other atom slightly positively charged so they attract
like magnets. This is an ionic bond.
Covalent bonds:
Ionic bonds:
Figure 2-3 Ionic Bonding
Section 2-1
Sodium atom (Na)
Chlorine atom (Cl)
Sodium ion (Na+)
Chloride ion (Cl-)
Transfer
of electron
Protons +11
Electrons -11
Charge
0
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Section:
Protons +17
Electrons -17
Charge
0
Protons +11
Electrons -10
Charge
+1
Protons +17
Electrons -18
Charge
-1
A structural formula shows which atoms bond where and by
sharing or “stealing” how many electrons. Ex:
Chemical formula: CO2
Structural formula: O = C = O This shows carbon is
between the two oxygen atoms and that 2 electrons are
shared between each oxygen and the carbon.
All atoms (except Helium and the very large atoms) want 8
electrons in their outer energy levels of electrons. This
stabilizes them. This is the reason why chemical bonds
form.
The first energy level is stable with 2 electrons (He only has
this energy level present). The next energy level needs
8 electrons but the first level must always be filled before
going to the next level. So an atom with 10 electrons is
stable (2 in first energy level and 8 in next.) An atom
with 18 electrons is also stable (2, 8, 8).
When combining substances, we can use a chemical
equation to show what atoms, etc. are present and how
many of each. The chemicals we are adding together are
called the “reactants.” The result of the combination is
called the “product.”
Ex: 2 Na + Cl2  2 NaCl
*Note: an arrow is used instead of an equal sign in chemical
equations. But, like in math, the number of each “variable”
(think of the elements added as “x’s” and “y’s”) must be
equal on each side of the equation.
The coefficient “2” in front of sodium (Na) tells us 2 separate
atoms of sodium were part of the reactants.
The subscript 2 on the Cl tells us a single molecule of
chlorine (containing 2 atoms of Cl) was the other reactant.
Note: a subscript only reflects back to the element
immediately in front of it unless parentheses are used.
2 Na + Cl2  2 NaCl
The product, 2 NaCl, tells us 2 molecules (EACH molecule
containing one Na and one Cl ) were produced.
Looking at other examples of chemical formulas using subscripts:
MgCl2
(1 magnesium, 2 chlorine atoms)
Mg(HCO3)2 (1 Mg atom, 2 molecules of bicarbonate {HCO3}
each containing 1 hydrogen, 1 carbon, and 3 oxygen
atoms) Since the 2 is directly outside the parentheses, it
distributes to everything within the parentheses. However,
the 3 is next to the “O” and no parentheses is in front of it
so it only applies to the oxygen. Total atoms: 1 Mg, 2 H, 2 C, 6 O
How many of each type of atom do I have if I write 6 Mg(HCO3)2?
6 Mg, 12 H, 12 C, 36 Oxygen
Figure 2-2 Isotopes of Carbon
Section 2-1
Nonradioactive carbon-12
Nonradioactive carbon-13
6 electrons
6 protons
6 neutrons
6 electrons
6 protons
7 neutrons
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Section:
Radioactive carbon-14
6 electrons
6 protons
8 neutrons
Isotopes are atoms that have the same number of protons and
electrons as the common atom of that type BUT have a different
number of neutrons. These isotopes are often unstable and
radioactive. Ex: Carbon 12 is stable with 6 neutrons (and 6
protons). The isotope, Carbon 14, has 8 neutrons and 6 protons. It
breaks down the extra neutrons over time (radioactivity). Carbon
dating measures the ratio of carbon 14 to carbon 12.
Ions are atoms that have become charged (usually in a solution) by
gaining or losing electrons. Ex: NaCl sometimes breaks into
separate Na+ and Cl- ions in solution. Chlorine gains an electron (it
already had 7 electrons in its outer energy level so it needed 1 to
complete the level) and becomes negatively charged. Since sodium
only had 1 electron in its outer energy level, it loses an electron to
become positively charged.
Even water can ionize. H2O molecules “dissociate’ to
become H3O+ cations (positively charged ions) and OHanions (negatively charged ions).
H3O+ is called a hydronium ion
OH- is called a hydroxide ion
When a solution has an equal number of hydronium (like
pure water) and hydroxide ions, it is neutral.
A solution with extra hydronium ions is acidic.
A solution with fewer hydronium ions than hydroxides is
basic, or alkaline.
Ex 1: HCl + H2O  H3O+ (aq) + Cl- (aq) (solution is acidic)
Ex 2: NaOH + H2O  Na+(aq) + OH-(aq) (solution is basic)
Acids are proton donors. Acids tend to taste sour (Ex:
lemon juice, vinegar)
Many chemical formulas for acids are written starting with
an “H” (the proton that is donated comes from hydrogen)
Ex: HCl, H2SO4, HCO3
Bases are proton acceptors. Bases tend to taste bitter.
Basic solutions are called alkaline. Strong bases are just as
caustic (corrosive, burning) as strong acids.
An acid reacted with a base can create a “salt” and water.
Ex: NaOH + HCl  NaCl + H2O
Acidity is measured on a pH scale. pH 7 = neutral,
pH 0 = highly acidic pH 14 = highly alkaline
Buffers help bring a solution closer to neutral.
Water
Water is VERY unique. Water is polar. That is, the oxygen atom in water
pulls harder on the electrons than the hydrogens do. So, the oxygen
“edge” is more negative than the hydrogen portions. This makes water
almost magnetic, so it likes to cling to surfaces. This is called adhesion.
Because water is polar, ionic substances, like NaCl, dissolve easily in it.
Nonpolar substances, like fat, don’t dissolve in water.
Water forms drops and bubbles up on surfaces because it likes to cling to
itself. This is called cohesion.
Water molecules form hydrogen bonds BETWEEN MOLECULES,
(Covalent and ionic bonds occur within a molecule or compound)
causing water molecules to form groups of molecules instead of
remaining as individuals.
Why is this important?
Because individual water molecules are lighter than air so water could not
exist as a liquid if it didn’t group together! (Air is mostly N2 gas)
As temperature drops to freezing, water sets up a crystalline structure that
spreads the molecules out a bit when forming ice. This makes water
LESS DENSE as a solid! It is the only Earthly substance that is less
dense as a solid.
Hydrogen bonds between water molecules (covalent bonds
within water molecules).
Mixtures, Solutions and Suspensions
A mixture - 2 or more elements or compounds physically mixed together but
not chemically combined. Ex: salt mixed with sand
When something like NaCl (salt) is mixed in water and the salt breaks into
Na+ and Cl- ions, they equally disperse. This is a solution. The salt and
water do not recombine to make a new substance such as NaOH and
HCl so this is still a mixture. The taste of salt is still distinct letting us
know it didn’t chemically react.
The water is called a solvent. (Because it is polar, water can dissolve ionic
and polar substances easily.) Solvents dissolve other things (Alcohol is
nonpolar so it can dissolve nonpolar substances. In chemistry, like
dissolves like) Concentration – amount of solute dissolved in solvent
The salt, or anything being dissolved into something else, is the solute.
Suspensions occur when molecules don’t dissolve by breaking into ions but
are small enough to remain suspended within the liquid if it is stirred or
moved often enough. If the mixing action stops for too long, the
substance will settle out. Ex: Blood cells in blood plasma. “Shake well”
Colloids, or colloidal suspensions, don’t settle out even if not stirred. Ex: milk
Organic Chemistry
Organic chemistry is also called the chemistry of carbon.
Carbon has 4 electrons in its outer energy layer. This means it needs
to either gain or lose 4 electrons to be happy. It pulls on its own
electrons hard enough that they can’t be stripped away yet it isn’t
strong enough to “steal” electrons from other atoms. So carbon
shares (covalently bonds) easily with other atoms, sometimes
sharing 1, 2, or 3 electrons with another atom. (In other words,
forming single, double, or triple bonds.)
Carbons can bond together to form chains, branched chains, or rings.
Carbon can form large molecules, called macromolecules, that are
essential to life. These include:
Carbohydrates (simple and complex sugars)
Lipids (fats, oils, waxes, steroids - think “cholesterol”)
Proteins (includes enzymes, hormones, etc.)
Nucleic acids (DNA, RNA)
Figure 2-11 Carbon Compounds
Section 2-3
Methane
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Section:
Acetylene
Butadiene
Benzene
Isooctane
Carbohydrates are made of carbon, hydrogen, and oxygen, usually in a ratio
of about 1:2:1 Ex: C6H12O6 (glucose)
Carbohydrates with fewer carbons are called simple sugars. These include
monosaccharides (the building blocks, or monomers, basic unit of
sugars) and disaccharides. Most simple sugars are given names that
end in “-ose.” Ex: glucose, fructose, lactose, dextrose, maltose, etc.
Complex sugars (polymers) like starch and glycogen are called
polysaccharides.
(polymers are more than one monomer joined together)
Simple sugars generally dissolve in water but complex sugars do not.
---------------------------------
Lipids are also made of carbon, hydrogen, and oxygen but have far fewer
oxygen atoms than carbohydrates. Fatty chains of carbon atoms with
many double bonds between carbons are called polyunsaturated fats.
Chains with all single bonds are called saturated fats. Polyunsaturated
fats are usually liquid at room temperature (Ex: corn oil) Saturated fats
like butter, animal fat, etc. are solids at room temperature.
Every cell in our body has a membrane made of lipids!
Figure 2-13 A Starch
Section 2-3
Starch
Glucose
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Section:
Proteins are made of C, H, O, and nitrogen that form amino acids, or
protein monomers (building blocks) (amine = NH2, carboxylic acid = -COOH)
Many different types of amino acids join via peptide bonds to form proteins.
Some amino acids are polar and some are nonpolar.
All proteins have a natural shape, or conformation. Some are folded,
some form complex groups of subunits, etc. This natural shape is
important to protein function. Changes in protein shape can damage or
destroy its ability to perform. Heat, radiation, etc. can denature (unfold
and destroy a protein’s function).
Hormones, muscles, most bodily structures, enzymes, etc. are made of
proteins.
Enzymes are special proteins that help a chemical reaction take place,
much like heat helps cake batter become an actual cake. The heat is
not an actual ingredient in the cake, it doesn’t chemically react with the
ingredients, yet without it, no cake. This is how enzymes behave in
chemical reactions. They lower the “activation energy” to get a reaction
going.
Enzymes that speed up reactions are called catalysts.
The reactants in an enzyme “catalyzed” reaction are called substrates.
Effect of Enzymes
Section 2-4
Reaction pathway
without enzyme
Activation energy
without enzyme
Reactants
Reaction pathway
with enzyme
Activation
energy
with enzyme
Products
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Section:
Figure 2-16 Amino Acids
Section 2-3
Amino group
Carboxyl group
General structure
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Section:
Alanine
Serine
Nucleic acids contain C, H, O, N, and phosphorus
monomers known as nucleotides. Each nucleotide
contains a 5 carbon sugar ring (ribose or deoxyribose), a
phosphate (PO4-) group, and a nitrogenous base
(adenine, guanine, cytosine, thymine, uracil).
DNA (deoxyribonucleic acid) is found in the nucleus of most
cells and controls their functions as well as determining
our inherited traits.
RNA (ribonucleic acid) helps carry out the DNA instructions
to make proteins, etc.
Interestingly, ATP (our cells’ energy source) is very similar
to RNA and DNA. It contains a 5 carbon sugar, the
nitrogenous base, adenine, and 3 phosphate groups.
Suggestion: Make a concept map for macromolecules.
Carbon Compounds
Macromolecules
include
Examples
Examples
Examples
Examples
which contain
which contain
which contain
which contain
Carbon,
hydrogen,
oxygen
Carbon,
hydrogen,
oxygen
Energy plays an important role in everything.
Potential energy is the “ability” to do work (movement of mass). It is
stored energy or energy due to position. (Ex: fat is stored energy, a
battery not in use has potential energy, a rock at the edge of a cliff
has potential energy, chemical bonds have potential energy)
Kinetic energy is energy of motion. (Current running through a wire is
kinetic energy, movement is kinetic energy.)
It is important to remember that kinetic energy is what makes a
substance a gas, liquid, or solid. The faster the molecules move in a
substance, the more kinetic energy it has.
Solids have molecules that are barely moving, just vibrating in place.
(Low kinetic energy)
Liquids have molecules bouncing off each other and sliding past one
another. (Medium kinetic energy)
Gases have molecules rapidly moving, slamming against each other
and bouncing far distances apart. (High kinetic energy)
Heating a substance can increase its kinetic energy.
Energy can be converted from one form to another. Ex:
Mechanical to electrical (generators, alternators);
chemical to heat & light (burning a match), solar to
chemical (photosynthesis)
Molecules can hold energy (potential energy) in their
bonds. Breaking or forming bonds during a chemical
reaction can consume (absorb) or release energy, often
in the form of heat.
Activation energy is the energy needed to get a reaction
started.
Reactions that give off more heat energy than they absorb
are called exothermic reactions.
Reactions that use more heat energy than they produce
are called endothermic reactions.