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Transcript
Chemistry
Matter
• Organisms are composed of matter
• Matter is anything that takes up space and has mass
• Matter is composed of chemical elements
• Matter is found on the Earth in three physical states
– Solid
– Liquid
– Gas
States of Matter
• Gases take the shape and volume of their container and can be compressed to
form liquids.
• Liquids take the shape of their container, but they do have their own volume
• Solids are rigid and have a definite shape and volume.
Classification of Matter
• Element: a substance composed of only one type of atom (all the
atoms have the same number of protons).
• Molecule: a unit composed of two or more atoms joined together by
chemical bonds
• Compound: a substance composed of 2 or more elements that have
been joined by chemical bonds
• Mixture: a combination of 2 or more substances that do NOT
chemically bond e.g. sugar mixed with salt
Sodium
Chlorine
Sodium chloride
Pure Substances and Mixtures
Elements
•
If a pure substance cannot be decomposed into something else, then the substance is an element
•
There are 114 elements known, 92 naturally occurring
•
Each element is given a unique chemical symbol (one or two letters)
Periodic Table
Essential Elements Of Life
• Only about 25 of the elements are essential to life
• Carbon, hydrogen, oxygen, and nitrogen make up 96% of living
matter
• Most of the remaining 4% consists of calcium, phosphorus,
potassium, and sulfur
• Trace elements are those required by an organism in minute
quantities
Essential Elements Of Life
Periodic Chart
Atoms
• Each element consists of one kind of unique atom
• An atom is the smallest unit of matter that still retains the
properties of an element, it cannot be broken down to other
substances by chemical reactions
Subatomic Particles
• Atoms are composed of subatomic
particles
Nucleus
• Relevant subatomic particles include:
– Neutrons (no electrical charge)
(a)
– Protons (positive charge)
(b)
– Electrons (negative charge)
2
• Neutrons and protons form the atomic
nucleus
• Electrons form a cloud around the
nucleus
Cloud of
negative
Protons charge (2
electrons)
2
Neutrons
2
Electrons
Atomic Number And Atomic Mass
• Atoms of the various elements differ in number of
subatomic particles
• An element’s atomic number is the number of protons
• The number of protons (atomic number) determines the
element’s properties
• An element’s mass number is the sum of protons plus
neutrons in the nucleus
• Atomic mass, the atom’s total mass, can be approximated by
the mass number
Periodic Chart
Atomic number
Element symbol
Mass number
Orbitals
•
Electrons orbit the nucleus of an atom in specific electron shells
•
Each Orbital holds a maximum of 2 electrons each
•
Several orbitals may be the same distance from the nucleus and thus contain electrons
of the same energy. Such electrons are said to occupy the same energy level or shell.
•
Rule of Eights for filling each shell:
First electron shell
(can hold 2 electrons)
Outermost electron shell
(can hold 8 electrons)
Electron
Hydrogen (H)
Atomic number = 1
Carbon (C)
Atomic number = 6
Nitrogen (N)
Atomic number = 7
Oxygen (O)
Atomic number = 8
Electron Shell Significance
• Electrons determine how an atom behaves
when it encounters other atoms
• Outer orbital (valence shell) determines
reactivity of atom - Electronegativity
• Atoms “desire” full outer orbitals
– Give up electrons (Na)
– Take electrons (Cl)
– Share electrons (O2)
• Noble gases - full outer shells (inert)
Chemical Bonding and Molecules
• Chemical reactions enable atoms to give up or
acquire electrons in order to complete their outer
shells
– These interactions usually result in atoms staying
close together
– The atoms are held together by chemical bonds
Kinetic Theory Of Matter
• All atoms and molecules are in constant random motion.
(Energy of motion is called kinetic energy.)
• The higher the temperature, the faster the atoms and
molecules move.
• All motion theoretically stops at absolute zero.
Energy
• Energy is the capacity to do work or ability to
cause change. Any change in the universe
requires energy. Energy comes in 2 forms:
– Potential energy is stored energy. No change is
currently taking place
– Kinetic energy is currently causing change.
This always involves some type of motion.
Forms Of Energy
• Kinetic energy is the energy
associated with motion
On the platform, a diver
has more potential energy.
Diving converts potential
energy to kinetic energy.
• Potential energy
– Is stored in the location of
matter
– Includes chemical energy
stored in molecular structure
• Energy can be converted from
one form to another
• First Law Of Thermodynamics
states that energy cannot be
created or destroyed; energy
can be transferred or
transformed
Climbing up converts kinetic
energy of muscle movement
to potential energy.
In the water, a diver has
less potential energy.
Temperature, Pressure, And Volume
• Volume – Pressure Relationship
– At a constant temperature, volume is
inversely proportional to pressure
• Volume – Temperature Relationship
– At constant pressure, the volume is
directly proportional to temperature
Chemical Reactions
• Cells constantly rearrange molecules by breaking
existing chemical bonds and forming new ones
• Such changes in the chemical composition of
matter are called chemical reactions
• Chemical reactions enable atoms to give up or
acquire electrons in order to complete their outer
shells
– These interactions usually result in atoms staying close
together
– The atoms are held together by chemical bonds
• Reactions can be written as equations
Chemical Equations
• The chemical equation for the formation of water
can be visualized as two hydrogen molecules
reacting with one oxygen molecule to form two
water molecules:
• 2H2 + O2  2H2O
Hydrogen gas Oxygen gas
Reactants
Reactants
Water
Products
Reading Chemical Equations
• The plus sign (+) means “react” and the arrow points towards the
substance produce in the reaction
• The chemical formulas on the right side of the equation are called reactants
and after the arrow are called product
• The numbers in front of the molecules or atoms indicate the number of
individual molecules or atoms (stoichiometric coefficients)
• The numbers behind are subscripts indicating the molecules or atoms are
bonded
2Na + 2H2O  2NaOH + H2
Reactants
Products
Chemical Reactions
• Are dependent on :
– Concentration
– Speed
– Energy (energy of activation)
– Orientation
Types Of Chemical Reactions
• Synthesis reactions - atoms or molecules combine to
form a product
• Decomposition reactions - molecules breakdown into
smaller molecules or atoms
• Exchange reactions - molecules exchange constituent
components (swap partners)
• Reversible reactions - the product of a previous
reaction can revert to the original reactants.
Combination (Synthesis) Reactions
• Combination (Synthesis) reaction occurs when
two or more substances react to form products:
Na + Cl  NaCl
Ca + 2NaCl  CaCl2
• In both cases, Sodium and Calcium combine
with Chlorine
Decomposition Reactions
• Decomposition reaction is when one substance
undergoes a reaction to produce two or more
substances:
2H2O  2H2 + O2
H2O2  H2O + O2
Exchange Reactions
• Exchange reaction occurs when molecules
“swap partners”:
NaOH + HCl  NaCl + H2O
H2CO3 + NaOH  H2O + NaHCO3
Reversible Reactions
• Reversible reactions can go forwards
(decomposition) or backwards (combination):
H2CO3  CO2 + H2O
• Chemical Equilibrium is defined as the state of
dynamic balance in which the rates of forward and
reverse processes (reactions) are equal
Chemical Products
• Element: a substance composed of only one type of atom
(all the atoms have the same number of protons).
• Molecule: a unit composed of two or more atoms joined
together by chemical bonds
• Compound: a substance composed of 2 or more elements
that have been joined by chemical bonds
• Mixture: a combination of 2 or more substances that do
NOT chemically bond e.g. sugar mixed with salt
Ionic Bonds
•
Atoms sometimes strip electrons from their bonding partners
•
An example is the transfer of an electron from sodium to chlorine
•
After the transfer of an electron, both atoms have charges
•
A charged atom (or molecule) is called an ion
•
–
An anion is a negatively charged ion
–
A cation is a positively charged ion
An ionic bond is an attraction between an anion and a cation - oppositely
charged ions
Ions And Ionic Compounds
• When an atom or molecule loses electrons, it becomes
positively charged.
– For example, when Na loses an electron it becomes Na+.
• Positively charged ions are called cations.
• When an atom or molecule gains electrons, it becomes
negatively charged.
• For example when Cl gains an electron it becomes Cl-.
• Negatively charged ions are called anions.
• An atom or molecule can lose more than one electron.
• When molecules loose electrons, polyatomic ions are
formed.
Ionic Compounds
• Compounds formed by ionic bonds are
called ionic compounds, or salts
• Salts, such as sodium chloride (table salt),
are often found in nature as crystals
Na+
Cl–
Covalent Bonds
• Molecules are formed by covalent bonds
Hydrogen atoms (2 H)
• A covalent bond is when two atoms share
one or more pairs of outer-shell electrons
(valence electrons)
• In a covalent bond, the shared electrons
count as part of each atom’s valence shell
• Much stronger than ionic bonds – holds lots
of Energy
• A single covalent bond, or single bond, is
the sharing of one pair of valence electrons
• A double covalent bond, or double bond, is
the sharing of two pairs of valence electrons
• Covalent bonds can form between atoms of
the same element or atoms of different
elements
Hydrogen
molecule (H2)
Covalent Bonds
Name
(molecular
formula)
Oxygen (O2)
Electronshell
diagram
Structural
formula
Spacefilling
model
Covalent Bonds
Name
(molecular
formula)
Electronshell
diagram
Structural
formula
Spacefilling
model
Water (H2O)
Name
(molecular
formula)
Methane (CH4)
Electronshell
diagram
Structural
formula
Spacefilling
model
Covalent Bonds
Figure 2.9
Electronegativity
• Outer orbital (valence shell) determines reactivity of atom Electronegativity
• Electronegativity is an atom’s attraction for the electrons in a covalent
bond
• The more electronegative an atom, the more strongly it pulls shared
electrons toward itself
–
O
+
H
H
H2O
+
Polar Covalent Bond
• In a nonpolar covalent bond, the atoms share the
electron equally
• In a polar covalent bond, one atom is more
electronegative, and the atoms do not share the
electron equally
The Structure Of Water
– Its two hydrogen atoms are joined to one oxygen atom by
single covalent bonds
– But the electrons of the covalent bonds are not shared equally
between oxygen and hydrogen
– This unequal sharing makes water a polar molecule
()
()
(-)
(-)
Unnumbered Figure 2.2
Hydrogen Bonds
• A hydrogen bond forms
when a hydrogen atom
covalently bonded to one
electronegative atom is also
attracted to another
electronegative atom
(-)
Hydrogen bond
()
()
• In living cells, the
electronegative partners are
usually oxygen or nitrogen
atoms
(-)
(-)
()
(-)
()
(b)
Figure 2.11b
Hydrogen Bonds
–
+
Water
(H2O)
+
Hydrogen bond
–
Ammonia
(NH3)
+
+
+
Weak Chemical Bonds
• Most of the strongest bonds in organisms are
covalent bonds that form a cell’s molecules
• Weak chemical bonds, such as ionic bonds and
hydrogen bonds, are also important
• Weak chemical bonds reinforce shapes of large
molecules and help molecules adhere to each
other
Biological Importance Of Water
• Acts as a powerful solvent
• Participates in chemical reactions
• Water has a high specific heat which moderates temperature - absorbs
and releases heat very slowly, minimizes temperature fluctuations to
within limits that permit life
– Heat is absorbed when hydrogen bonds break
– Heat is released when hydrogen bonds form
• Requires a great amount of heat to change to a gas
– Heat of vaporization - the quantity of heat a liquid must absorb for 1
gram of it to be converted from a liquid to a gas
– Evaporative cooling - Allows water to cool a surface due to water’s high
heat of vaporization
• Acts as a lubricant
Polarity & Hydrogen Bonds
• Cohesion - molecules attract other
water molecules
• Capillarity
– Water molecules are drawn up a narrow
tube
– Helps pull water up through the
microscopic vessels of plants
• Surface tension
– water molecules on the surface cling to
each other – related to cohesion
– Is a measure of how hard it is to break
the surface of a liquid
• Adhesion - water molecules attract
other charged substances
Water As A Solvent
• Water is a versatile solvent due to its polarity
• It can form aqueous solutions
• The different regions of the polar water molecule can interact with
ionic compounds called solutes and dissolve them
Negative
oxygen regions
of polar water molecules
are attracted to sodium
cations (Na+).
–
Na
+
+
–
Positive
hydrogen regions
of water molecules
cling to chloride anions
(Cl–).
–
Na
+
Cl–
+
Cl –
–
+
+
–
–
+
+
–
–
+
+
–
–
Formula and Molecular Weights
• Formula weights (FW) is the sum of the atomic weights of
each atom in the chemical formula.
• FW (H2SO4) = 2AW(H) + AW(S) + 4AW(O)
• = 2(1.0 amu) + (32.0 amu) + 4(16.0)
• = 98.0 amu
• If the chemical formula is also its molecular formula then
the weight is called the molecular weight (MW).
• MW(C6H12O6) = 6(12.0 amu) + 12(1.0 amu) + 6(16.0
amu) =????
The Mole
• The unit we use to express the quantity of atoms, ions, and
molecules that an object contains is called mole.
– Mole: convenient measure chemical quantities.
• The actual number of atoms, ions, or molecules in 1 mole of
something = 6.0221367  1023 (Advogadro’s number).
• Thus,
• 1 mole of 12C atoms = 6.02 x 1023 12C atoms
• 1 mole of H2O molecules = 6.02 x 1023 molecules
• 1 mole of NO3- ions = 6.02 x 1023 ions
Visualizing The Mole Concept
Different Units
Solution Composition
• Solutions are homogenous mixtures of two or more substances:
– Solute: present in smallest amount and is the substance dissolved in the
solvent.
– Solvent: present in the greater quantities and is used to dissolve the solute.
– Example: NaCl dissolved in Water (water = Solvent and NaCl = Solute)
• Change concentration by using different amounts of solute and solvent.
• Molarity: Moles of solute per liter of solution.
Concentration of Solutions
• Percent solutions
• Ratio of solute to solvent expressed as a
percentage: weight (g)/volume (ml)
• Unit seen on IV bags and medicinal solutions
– 5% dextrose = 5g dextrose / 100 ml of solution
– 0.9% saline = 0.9g NaCl / 100 ml of solution
Example
• Betadine antiseptic solution contains 10g of
povidine-iodine in 100mL of solution. Calculate
the percent (w/v) concentration of the solution.
% w/v =
grams of solute
x 100mL of solution
= 10g/100mL x 100
=
10%
Concentrations Of Solutions
Formula for Molarity
• The most widely used way of quantifying concentration
of solutions in chemistry. Molarity is generally
represented by the symbol M and defined as the number
of moles of solute dissolved in a liter of solution.
General Properties Of Aqueous
Solutions (Terms)
• Acids - substances that able to ionize in solution to form
hydrogen ion (H+) and increase the concentration of H+ in
the solution.
• For example, HCl dissociate in water to form H+ and Clions.
• Bases - are substances that can react with or accept H+ ions.
• For example, OH- will accept H+ from HCl forming H2O.
• Salts - are ionic compounds that can be formed by replacing
one or more of the hydrogen ions of an acid by a different
positive ion.
• For example, NaCl instead of HCl.
Acids, Bases, And pH
• Dissociation of water molecules leads to acidic and basic
conditions that affect living organisms
– Water dissociates into hydronium ions and hydroxide ions
– Changes in the concentration of these ions can have a great
affect on living organisms
–
+
H
H
H
H
H
H
H
Hydronium
ion (H3O+)
+
H
Hydroxide
ion (OH–)
Acids, Bases, And pH
• Acid - A chemical compound that dissociates into
one or more hydrogen ions (H+) and one or more
negative ions (anions). An acid donates H+ ions
(protons) to solutions
• Base - Dissociates into one or more positive ions
(cations) and one or more hydroxide ions (OH-). A
base accepts H+ ions and removes them from
solution, reducing the hydrogen ion concentration
of a solution
Strong And Weak Acids
• Strong Acids ionize or dissociate completely in
water
HCl + H2O  H+(aq) + Cl-(aq)
– Strong Acids: HNO3 (nitric acid); H2SO4 (sulfuric
acid)
• Weak Acids ionize slightly (less than 5%) in water
– CH3COOH (acetic acid)
– H2CO3 (carbonic acid)
– H3PO4 (phosphoric acid)
pH
• To describe the acidity of a
solution, we use the pH scale
• Is a measure of the concentration
of H+ ions in a solution
Oven cleaner
Household bleach
Household ammonia
Basic
solution
• Is determined by the relative
concentrations of H+
– Basic = High pH = few
many OH-
Milk of magnesia
Seawater
Human blood
Pure water
H +,
Neutral
solution
– Acidic = Low pH = many H+,
few OH-
Urine
Tomato juice
Grapefruit juice
Acidic
solution
Lemon juice;
gastric juice
pH scale
Figure 2.17
pH Scale
• The pH scale is a logarithmic scale
used to express the amount of H+
ions in a solution.
• pH is defined as the negative log of
the [H+] of a solution
• A change of one whole number
represents a tenfold (10X) change in
the number of H+ ions.
• A solution with pH 3 has ten times
as many H+ ions as a solution with
pH 4
Buffers
•
A buffer is a substance that helps minimize the change in the pH of a
solution when acids or bases are added.
•
Consist of an acid-base pair that reversibly combines with hydrogen
ions
•
Buffers work by releasing H+ when their concentration falls, and
absorbing H+ when their concentration rises.
•
Buffers are important to living organisms because most cells can survive
and function normally only within a relatively narrow range
1
-
pH
14
[H+]
Bicarbonate Buffer System
•
A mixture of carbonic acid (H2CO3) and its salt, sodium bicarbonate (NaHCO3)
•
If strong acid is added:
– Hydrogen ions released combine with the bicarbonate ions and form carbonic acid (a weak
acid)
– The pH of the solution decreases only slightly
•
If strong base is added:
– It reacts with the carbonic acid to form sodium bicarbonate (a weak base)
– The pH of the solution rises only slightly