PowerPoint® Lecture Slides prepared by Janice Meeking, Mount Royal College
C H A P T E R
Copyright © 2010 Pearson Education, Inc.
2
Chemistry Comes Alive: Part A
Copyright © 2010 Pearson Education, Inc.
Matter
• Anything that has mass and occupies space
• States of matter:
1. Solid—definite shape and volume
2. Liquid—definite volume, changeable shape
3. Gas—changeable shape and volume
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Energy
• Capacity to do work or put matter into motion
• Types of energy:
• Kinetic—energy in action
• Potential—stored (inactive) energy
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Forms of Energy
• Chemical energy—stored in bonds of chemical substances
• Electrical energy—results from movement of charged particles
• Mechanical energy—directly involved in moving matter
• Radiant or electromagnetic energy—exhibits wavelike properties (i.e., visible light, ultraviolet light, and X-rays)
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Energy Form Conversions
• Energy may be converted from one form to another
• Conversion is inefficient because some energy is “lost” as heat
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Composition of Matter
• Elements
• Cannot be broken down by ordinary chemical means
• Each has unique properties:
• Physical properties
• Are detectable with our senses, or are measurable
• Chemical properties
• How atoms interact (bond) with one another
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Composition of Matter
• Atoms
• Unique building blocks for each element
• Atomic symbol: one- or two-letter chemical shorthand for each element
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Major Elements of the Human Body
• Oxygen (O)
• Carbon (C)
• Hydrogen (H)
• Nitrogen (N)
About 96% of body mass
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Lesser Elements of the Human Body
• About 3.9% of body mass:
• Calcium (Ca), phosphorus (P), potassium (K), sulfur (S), sodium (Na), chlorine (Cl), magnesium (Mg), iodine (I), and iron (Fe)
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Trace Elements of the Human Body
• < 0.01% of body mass:
• Part of enzymes, e.g., chromium (Cr), manganese (Mn), and zinc (Zn)
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Atomic Structure
• Determined by numbers of subatomic particles
• Nucleus consists of neutrons and protons
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Atomic Structure
• Neutrons
• No charge
• Mass = 1 atomic mass unit (amu)
• Protons
• Positive charge
• Mass = 1 amu
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Atomic Structure
• Electrons
• Orbit nucleus
• Equal in number to protons in atom
• Negative charge
• 1/2000 the mass of a proton (0 amu)
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Models of the Atom
• Orbital model: current model used by chemists
• Depicts probable regions of greatest electron density (an electron cloud)
• Useful for predicting chemical behavior of atoms
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Models of the Atom
• Planetary model—oversimplified, outdated model
• Incorrectly depicts fixed circular electron paths
• Useful for illustrations (as in the text)
Copyright © 2010 Pearson Education, Inc. Figure 2.1
(a) Planetary model (b) Orbital model
Helium atom
2 protons (p+)2 neutrons (n0)2 electrons (e–)
Helium atom
2 protons (p+)2 neutrons (n0)2 electrons (e–)
Nucleus Nucleus
Proton Neutron Electroncloud
Electron
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Identifying Elements
• Atoms of different elements contain different numbers of subatomic particles
• Compare hydrogen, helium and lithium (next slide)
Copyright © 2010 Pearson Education, Inc. Figure 2.2
Proton
Neutron
Electron
Helium (He)(2p+; 2n0; 2e–)
Lithium (Li)(3p+; 4n0; 3e–)
Hydrogen (H)(1p+; 0n0; 1e–)
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Identifying Elements
• Atomic number = number of protons in nucleus
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Identifying Elements
• Mass number = mass of the protons and neutrons
• Mass numbers of atoms of an element are not all identical
• Isotopes are structural variations of elements that differ in the number of neutrons they contain
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Identifying Elements
• Atomic weight = average of mass numbers of all isotopes
Copyright © 2010 Pearson Education, Inc. Figure 2.3
Proton
Neutron
Electron
Deuterium (2H)(1p+; 1n0; 1e–)
Tritium (3H)(1p+; 2n0; 1e–)
Hydrogen (1H)(1p+; 0n0; 1e–)
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Radioisotopes
• Spontaneous decay (radioactivity)
• Similar chemistry to stable isotopes
• Can be detected with scanners
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Radioisotopes
• Valuable tools for biological research and medicine
• Cause damage to living tissue:
• Useful against localized cancers
• Radon from uranium decay causes lung cancer
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Molecules and Compounds
• Most atoms combine chemically with other atoms to form molecules and compounds
• Molecule—two or more atoms bonded together (e.g., H2 or C6H12O6)
• Compound—two or more different kinds of atoms bonded together (e.g., C6H12O6)
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Mixtures
• Most matter exists as mixtures
• Two or more components physically intermixed
• Three types of mixtures
• Solutions
• Colloids
• Suspensions
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Solutions
• Homogeneous mixtures
• Usually transparent, e.g., atmospheric air or seawater
• Solvent
• Present in greatest amount, usually a liquid
• Solute(s)
• Present in smaller amounts
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Concentration of Solutions
• Expressed as
• Percent, or parts per 100 parts
• Milligrams per deciliter (mg/dl)
• Molarity, or moles per liter (M)
• 1 mole = the atomic weight of an element or molecular weight (sum of atomic weights) of a compound in grams
• 1 mole of any substance contains 6.02 1023 molecules (Avogadro’s number)
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Colloids and Suspensions
• Colloids (emulsions)
• Heterogeneous translucent mixtures, e.g., cytosol
• Large solute particles that do not settle out
• Undergo sol-gel transformations
• Suspensions:
• Heterogeneous mixtures, e.g., blood
• Large visible solutes tend to settle out
Copyright © 2010 Pearson Education, Inc. Figure 2.4
Solution
Soluteparticles
Soluteparticles
Soluteparticles
Solute particles are verytiny, do not settle out or
scatter light.
ColloidSolute particles are larger
than in a solution and scatterlight; do not settle out.
SuspensionSolute particles are very
large, settle out, and mayscatter light.
ExampleMineral water
ExampleGelatin
ExampleBlood
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Mixtures vs. Compounds
• Mixtures
• No chemical bonding between components
• Can be separated physically, such as by straining or filtering
• Heterogeneous or homogeneous
• Compounds
• Can be separated only by breaking bonds
• All are homogeneous
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Chemical Bonds
• Electrons occupy up to seven electron shells (energy levels) around nucleus
• Octet rule: Except for the first shell which is full with two electrons, atoms interact in a manner to have eight electrons in their outermost energy level (valence shell)
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Chemically Inert Elements
• Stable and unreactive
• Outermost energy level fully occupied or contains eight electrons
Copyright © 2010 Pearson Education, Inc. Figure 2.5a
Helium (He)(2p+; 2n0; 2e–)
Neon (Ne)(10p+; 10n0; 10e–)
2e 2e8e
(a) Chemically inert elements
Outermost energy level (valence shell) complete
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Chemically Reactive Elements
• Outermost energy level not fully occupied by electrons
• Tend to gain, lose, or share electrons (form bonds) with other atoms to achieve stability
Copyright © 2010 Pearson Education, Inc. Figure 2.5b
2e4e
2e8e
1e
(b) Chemically reactive elementsOutermost energy level (valence shell) incomplete
Hydrogen (H)(1p+; 0n0; 1e–)
Carbon (C)(6p+; 6n0; 6e–)
1e
Oxygen (O)(8p+; 8n0; 8e–) Sodium (Na)
(11p+; 12n0; 11e–)
2e6e
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Ionic Bonds
• Ions are formed by transfer of valence shell electrons between atoms
• Anions (– charge) have gained one or more electrons
• Cations (+ charge) have lost one or more electrons
• Attraction of opposite charges results in an ionic bond
Copyright © 2010 Pearson Education, Inc. Figure 2.6a-b
Sodium atom (Na)(11p+; 12n0; 11e–)
Chlorine atom (Cl)(17p+; 18n0; 17e–)
Sodium ion (Na+) Chloride ion (Cl–)
Sodium chloride (NaCl)
+ –
(a) Sodium gains stability by losing one electron, and chlorine becomes stable by gaining one electron.
(b) After electron transfer, the oppositely charged ions formed attract each other.
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Formation of an Ionic Bond
• Ionic compounds form crystals instead of individual molecules
• NaCl (sodium chloride)
Copyright © 2010 Pearson Education, Inc. Figure 2.6c
CI–
Na+
(c) Large numbers of Na+ and Cl– ions associate to form salt (NaCl) crystals.
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Covalent Bonds
• Formed by sharing of two or more valence shell electrons
• Allows each atom to fill its valence shell at least part of the time
Copyright © 2010 Pearson Education, Inc. Figure 2.7a
+
Hydrogenatoms
Carbonatom
Molecule ofmethane gas (CH4)
Structuralformulashows singlebonds.
(a) Formation of four single covalent bonds: carbon shares four electron pairs with four hydrogen atoms.
or
Resulting moleculesReacting atoms
Copyright © 2010 Pearson Education, Inc. Figure 2.7b
or
Oxygenatom
Oxygenatom
Molecule ofoxygen gas (O2)
Structuralformulashowsdouble bond.(b) Formation of a double covalent bond: Two
oxygen atoms share two electron pairs.
Resulting moleculesReacting atoms
+
Copyright © 2010 Pearson Education, Inc. Figure 2.7c
+ or
Nitrogenatom
Nitrogenatom
Molecule ofnitrogen gas (N2)
Structuralformulashowstriple bond.(c) Formation of a triple covalent bond: Two
nitrogen atoms share three electron pairs.
Resulting moleculesReacting atoms
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Covalent Bonds
• Sharing of electrons may be equal or unequal
• Equal sharing produces electrically balanced nonpolar molecules
• CO2
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Covalent Bonds
• Unequal sharing by atoms with different electron-attracting abilities produces polar molecules
• H2O
• Atoms with six or seven valence shell electrons are electronegative, e.g., oxygen
• Atoms with one or two valence shell electrons are electropositive, e.g., sodium
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Hydrogen Bonds
• Attractive force between electropositive hydrogen of one molecule and an electronegative atom of another molecule
• Common between dipoles such as water
• Also act as intramolecular bonds, holding a large molecule in a three-dimensional shape
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(a) The slightly positive ends (+) of the watermolecules become aligned with the slightlynegative ends (–) of other water molecules.
+
–
–
–– –
+
+
+
+
+
Hydrogen bond(indicated bydotted line)
Figure 2.10a
Copyright © 2010 Pearson Education, Inc. Figure 2.10b
(b) A water strider can walk on a pond because of the highsurface tension of water, a result of the combinedstrength of its hydrogen bonds.
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Chemical Reactions
• Occur when chemical bonds are formed, rearranged, or broken
• Represented as chemical equations
• Chemical equations contain:
• Molecular formula for each reactant and product
• Relative amounts of reactants and products, which should balance
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Examples of Chemical Equations
H + H H2 (hydrogen gas)
4H + C CH4 (methane)
(reactants) (product)
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Patterns of Chemical Reactions
• Synthesis (combination) reactions
• Decomposition reactions
• Exchange reactions
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Synthesis Reactions
• A + B AB
• Always involve bond formation
• Anabolic
Copyright © 2010 Pearson Education, Inc. Figure 2.11a
ExampleAmino acids are joined together toform a protein molecule.
(a) Synthesis reactions
Smaller particles are bondedtogether to form larger,
more complex molecules.
Amino acidmolecules
Proteinmolecule
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Decomposition Reactions
• AB A + B
• Reverse synthesis reactions
• Involve breaking of bonds
• Catabolic
Copyright © 2010 Pearson Education, Inc. Figure 2.11b
ExampleGlycogen is broken down to releaseglucose units.
Bonds are broken in largermolecules, resulting in smaller,
less complex molecules.
(b) Decomposition reactions
Glucosemolecules
Glycogen
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Exchange Reactions
• AB + C AC + B
• Also called displacement reactions
• Bonds are both made and broken
Copyright © 2010 Pearson Education, Inc. Figure 2.11c
ExampleATP transfers its terminal phosphategroup to glucose to form glucose-phosphate.
Bonds are both made and broken(also called displacement reactions).
(c) Exchange reactions
Glucose Adenosine triphosphate (ATP)
Adenosine diphosphate (ADP)Glucosephosphate
+
+
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Oxidation-Reduction (Redox) Reactions
• Decomposition reactions: Reactions in which fuel is broken down for energy
• Also called exchange reactions because electrons are exchanged or shared differently
• Electron donors lose electrons and are oxidized
• Electron acceptors receive electrons and become reduced
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Chemical Reactions
• All chemical reactions are either exergonic or endergonic
• Exergonic reactions—release energy
• Catabolic reactions
• Endergonic reactions—products contain more potential energy than did reactants
• Anabolic reactions
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Chemical Reactions
• All chemical reactions are theoretically reversible
• A + B AB
• AB A + B
• Chemical equilibrium occurs if neither a forward nor reverse reaction is dominant
• Many biological reactions are essentially irreversible due to
• Energy requirements
• Removal of products
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