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2. Chemistry Comes Alive: Part A. Objectives: Chapter Two. Differentiate between matter and energy and between potential energy and kinetic energy. Define chemical element and list the four elements that form the bulk of body matter.

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slide1

2

Chemistry Comes Alive: Part A

objectives chapter two
Objectives: Chapter Two
  • Differentiate between matter and energy and between potential energy and kinetic energy.
  • Define chemical element and list the four elements that form the bulk of body matter.
  • Define molecule, and distinguish between a compound and a molecule.
  • Differentiate among ionic, covalent, and hydrogen bonds.
  • Compare and contrast polar and nonpolar bonds
  • Define the three major types of chemical reactions: synthesis, decomposition, and exchange. Comment on the nature of oxidation-reduction reactions and their importance.
  • Explain the importance of water and salts to body homeostasis.
objectives continued
Objectives Continued
  • Define acid and base, and explain the concept of pH.
  • Describe and compare the building blocks, general structures, and biological functions of carbohydrates and lipids.
  • Explain the role of dehydration synthesis and hydrolysis in the formation and breakdown of organic molecules.
  • Describe how enzymes function.
  • Explain the role of ATP in cell metabolism.
matter
Matter
  • Anything that has mass and occupies space
  • States of matter:
    • Solid—definite shape and volume
    • Liquid—definite volume, changeable shape
    • Gas—changeable shape and volume
energy
Energy
  • Capacity to do work or put matter into motion
  • Types of energy:
    • Kinetic—energy in action
    • Potential—stored (inactive) energy
forms of energy
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)
energy form conversions
Energy Form Conversions
  • Energy may be converted from one form to another, but is considered “inefficient”
  • Conversion is inefficient because some energy is given off as heat and may not be used. Heat is the lowest and most disorganized form of energy.
composition of matter
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
composition of matter1
Composition of Matter
  • Atoms
    • Unique building blocks for each element
  • Atomic symbol: one- or two-letter chemical shorthand for each element
major elements of the human body
Major Elements of the Human Body
  • Oxygen (O)
  • Carbon (C)
  • Hydrogen (H)
  • Nitrogen (N)

About 96% of body mass

lesser elements of the human body
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)
trace elements of the human body
Trace Elements of the Human Body
  • < 0.01% of body mass:
    • Part of enzymes, e.g., chromium (Cr), manganese (Mn), and zinc (Zn)
atomic structure
Atomic Structure
  • Determined by numbers of subatomic particles
  • Nucleus consists of neutrons and protons
atomic structure1
Atomic Structure
  • Neutrons
      • No charge
      • Mass = 1 atomic mass unit (amu)
  • Protons
      • Positive charge
      • Mass = 1 amu
atomic structure2
Atomic Structure
  • Electrons
    • Orbit nucleus
    • Equal in number to protons in atom
    • Negative charge
    • 1/2000 the mass of a proton (0 amu)
identifying elements
Identifying Elements
  • Atoms of different elements contain different numbers of subatomic particles
identifying elements1
Identifying Elements
  • Atomic number = number of protons in nucleus
identifying elements2
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
identifying elements3
Identifying Elements
  • Atomic weight = average of mass numbers of all isotopes
slide20

Proton

Neutron

Electron

Hydrogen (1H)

(1p+; 0n0; 1e–)

Deuterium (2H)

(1p+; 1n0; 1e–)

Tritium (3H)

(1p+; 2n0; 1e–)

Figure 2.3

radioisotopes
Radioisotopes
  • Spontaneous decay (radioactivity)
  • Similar chemistry to stable isotopes
  • Can be detected with scanners
radioisotopes1
Radioisotopes
  • Valuable tools for biological research and medicine
  • Cause damage to living tissue:
    • Useful against localized cancers
    • Radon from uranium decay causes lung cancer
molecules and compounds
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)
mixtures
Mixtures
  • Most matter exists as mixtures
    • Two or more components physically intermixed
  • Three types of mixtures
    • Solutions
    • Colloids
    • Suspensions
solutions
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
concentration of solutions
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)
mixtures vs compounds
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
chemical bonds
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)
chemically inert elements
Chemically Inert Elements
  • Stable and unreactive
  • Outermost energy level fully occupied or contains eight electrons
slide30

(a) Chemically inert elements

Outermost energy level (valence shell) complete

8e

2e

2e

Helium (He)

(2p+; 2n0; 2e–)

Neon (Ne)

(10p+; 10n0; 10e–)

Figure 2.5a

chemically reactive elements
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
slide32

(b) Chemically reactive elements

Outermost energy level (valence shell) incomplete

4e

2e

1e

Hydrogen (H)

(1p+; 0n0; 1e–)

Carbon (C)

(6p+; 6n0; 6e–)

1e

6e

8e

2e

2e

Oxygen (O)

(8p+; 8n0; 8e–)

Sodium (Na)

(11p+; 12n0; 11e–)

Figure 2.5b

types of chemical bonds
Types of Chemical Bonds
  • Ionic
  • Covalent
  • Hydrogen
ionic bonds
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
slide35

+

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.

Figure 2.6a-b

formation of an ionic bond
Formation of an Ionic Bond
  • Ionic compounds form crystals instead of individual molecules
    • NaCl (sodium chloride)
slide37

CI–

Na+

(c) Large numbers of Na+ and Cl– ions

associate to form salt (NaCl) crystals.

Figure 2.6c

covalent bonds
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
slide39

Reacting atoms

Resulting molecules

+

or

Structural

formula

shows

single

bonds.

Molecule of

methane gas (CH4)

Hydrogen

atoms

Carbon

atom

(a) Formation of four single covalent bonds:

carbon shares four electron pairs with four

hydrogen atoms.

Figure 2.7a

slide40

Reacting atoms

Resulting molecules

+

or

Structural

formula

shows

double

bond.

Molecule of

oxygen gas (O2)

Oxygen

atom

Oxygen

atom

(b) Formation of a double covalent bond: Two

oxygen atoms share two electron pairs.

Figure 2.7b

slide41

Reacting atoms

Resulting molecules

+

or

Structural

formula

shows

triple

bond.

Molecule of

nitrogen gas (N2)

Nitrogen

atom

Nitrogen

atom

(c) Formation of a triple covalent bond: Two

nitrogen atoms share three electron pairs.

Figure 2.7c

covalent bonds1
Covalent Bonds
  • Sharing of electrons may be equal or unequal
    • Equal sharing produces electrically balanced nonpolar molecules
      • CO2
covalent bonds2
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
hydrogen bonds
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

PLAY

slide48

+

–

Hydrogen bond

(indicated by

dotted line)

+

+

–

–

–

+

+

+

–

(a) The slightly positive ends (+) of the watermolecules become aligned with the slightlynegative ends (–) of other water molecules.

Figure 2.10a

slide49

(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.

Figure 2.10b

chemical reactions
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
examples of chemical equations
Examples of Chemical Equations

H + H  H2 (hydrogen gas)

4H + C  CH4 (methane)

(reactants)

(product)

patterns of chemical reactions
Patterns of Chemical Reactions
  • Synthesis (combination) reactions
  • Decomposition reactions
  • Exchange reactions
synthesis reactions
Synthesis Reactions
  • A + B  AB
    • Always involve bond formation
    • Anabolic
slide54

(a) Synthesis reactions

Smaller particles are bonded

together to form larger,

more complex molecules.

Example

Amino acids are joined together to

form a protein molecule.

Amino acid

molecules

Protein

molecule

Figure 2.11a

decomposition reactions
Decomposition Reactions
  • AB  A + B
    • Reverse synthesis reactions
    • Involve breaking of bonds
    • Catabolic
slide56

(b) Decomposition reactions

Bonds are broken in larger

molecules, resulting in smaller,

less complex molecules.

Example

Glycogen is broken down to release

glucose units.

Glycogen

Glucose

molecules

Figure 2.11b

exchange reactions
Exchange Reactions
  • AB + C  AC + B
    • Also called displacement reactions
    • Bonds are both made and broken
slide58

(c) Exchange reactions

Bonds are both made and broken

(also called displacement reactions).

Example

ATP transfers its terminal phosphate

group to glucose to form glucose-phosphate.

+

Glucose

Adenosine triphosphate (ATP)

+

Glucose

phosphate

Adenosine diphosphate (ADP)

Figure 2.11c

oxidation reduction redox reactions
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
chemical reactions1
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
chemical reactions2
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
rate of chemical reactions
Rate of Chemical Reactions
  • Rate of reaction is influenced by:
    •  temperature  rate
    •  particle size  rate
    •  concentration of reactant  rate
  • Catalysts:  rate without being chemically changed
    • Enzymes are biological catalysts