Chemistry Comes Alive - Dr. Jerry Cronindrjerrycronin.weebly.com/uploads/5/9/7/4/5974564/... · 2.3 Combining Matter Molecules and Compounds •Most atoms chemically combine with

PowerPoint® Lecture Slides

prepared by

Karen Dunbar Kareiva

Ivy Tech Community College© Annie Leibovitz/Contact Press Images

Chapter 2 Part A

Chemistry

Comes Alive

© 2017 Pearson Education, Inc.

Why This Matters:

• Understanding chemistry and biochemistry

helps to determine the most effective solutions

to use to treat dehydration and fluid loss



Video: Why This Matters

Chemistry and Physiological Reactions

• Body is made up of many chemicals

• Chemistry underlies all physiological reactions:

– Movement, digestion, pumping of heart, nervous

system

• Chemistry can be broken down into:

– Basic chemistry

– Biochemistry


Part 1 – Basic Chemistry

Matter

• Matter is anything that has mass and occupies

space

– Matter can be seen, smelled, and/or felt

– Weight is mass plus the effects of gravity


2.1 Matter and Energy

Matter

• States of matter

– Matter can exist in three possible states:

• Solid: definite shape and volume

• Liquid: changeable shape; definite volume

• Gas: changeable shape and volume


Energy

• Energy is the capacity to do work or put matter

into motion

• Energy does not have mass, nor does it take up

space

• The greater the work done, the more energy it

uses up


Animation – Energy Concepts


Energy (cont.)

• Kinetic versus potential energy

– Energy exists in two possible forms:

• Kinetic – energy in action

• Potential – stored (inactive) energy

– Energy can be transformed from potential to

kinetic energy

• Stored energy can be released, resulting in action


Energy (cont.)

• 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

• Travels in waves (example: heat, visible light,

ultraviolet light, and X rays)


Energy (cont.)

• Energy form conversions

– Energy may be converted from one form to

another

• Example: turning on a lamp converts electrical energy

to light energy

– Energy conversion is inefficient

• Some energy is “lost” as heat, which can be partly

unusable energy


2.2 Atoms and Elements

• All matter is composed of elements

– Elements are substances that cannot be broken

down into simpler substances by ordinary

chemical methods

• Four elements make up 96% of body:

– Carbon, oxygen, hydrogen, and nitrogen

– 9 elements make up 3.9% of body

– 11 elements make up <0.01%

• Periodic table lists all known elements



• All elements are made up of atoms, which are:

– Unique building blocks for each element

– Smallest particles of an element with properties

of that element

– What give each element its particular physical &

chemical properties



• Atomic symbol

– One- or two-letter chemical shorthand for each

element

• Example: “O” for oxygen, “C” for carbon

• Some symbols come from Latin names: “Na” (natrium)

is sodium; “K” (kalium) is potassium


Table 2.1-1 Common Elements Composing the Human Body


Table 2.1-2 Common Elements Composing the Human Body (continued)


Table 2.1-3 Common Elements Composing the Human Body (continued)


Structure of Atoms

• Atoms are composed of three subatomic

particles:

– Protons

• Carry a positive charge (+)

• Weigh an arbitrary 1 atomic mass unit (1 amu)

– Neutrons

• Have no electrical charge (0)

• Also weigh 1 amu

– Electrons

• Carry a negative charge (−)

• Are so tiny they have virtually no weight (0 amu)


Structure of Atoms (cont.)

• Number of positive protons is balanced by

number of negative electrons, so atoms are

electrically neutral

• Protons and neutrons are found in a centrally

located nucleus; electrons orbit around the

nucleus

• Chemists devise models of how subatomic

particles are put together

– Planetary model

– Orbital model


Structure of Atoms (cont.)

• Planetary model: simplified and outdated

because it incorrectly depicts electrons in orbits,

fixed circular paths

– Still useful for illustrations

• Orbital model: current model used that depicts

orbitals, probable regions where an electron is

most likely to be located (rather than fixed

orbits)

– Shading in regions of greatest electron density

results in an electron cloud around nucleus

– Useful for predicting chemical behavior of atoms© 2017 Pearson Education, Inc.

Figure 2.1 Two models of the structure of an atom.


Planetary model 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

Electron

cloud

Proton

Neutron

Electron

Identifying Elements

• Different elements contain different numbers of

subatomic particles

– Hydrogen has 1 proton, 0 neutrons, and 1

electron

– Helium has 2 protons, 2 neutrons, and 2

electrons

– Lithium has 3 protons, 4 neutrons, and 3

electrons

• Identifying facts about an element include its

atomic number, mass number, isotopes, and

atomic weight© 2017 Pearson Education, Inc.

Figure 2.2 Atomic structure of the three smallest atoms.


Proton

Neutron

Electron

Hydrogen (H)

(1p+; 0n0; 1e−)

Helium (He)

(2p+; 2n0; 2e−)

Lithium (Li)

(3p+; 4n0; 3e−)

Identifying Elements (cont.)

• Atomic number

– Number of protons in nucleus

– Written as subscript to left of atomic symbol

• Example: 3Li

• Mass number

– Total number of protons and neutrons in nucleus

• Total mass of atom

– Written as superscript to left of atomic symbol

• Example: 7Li


Identifying Elements (cont.)

• Isotopes

– Structural variations of same element

– Atoms contain same number of protons but differ

in the number of neutrons they contain

• Atomic numbers are same, but mass numbers

different

• Atomic weight

– Average of mass numbers of all isotope forms of

an atom


Figure 2.3 Isotopes of hydrogen.


Proton

Neutron

Electron

Hydrogen (1H)

(1p+; 0n0; 1e−)

Deuterium (2H)

(1p+; 1n0; 1e−)

Tritium (3H)

(1p+; 2n0; 1e−)

Radioisotopes

• Radioisotopes are isotopes that decompose to

more stable forms

– Atom loses various subatomic particles

• Sometimes loss results in an isotope becoming a

different element

– As isotope decays, subatomic particles that are

being given off release a little energy

• This energy is referred to as radioactivity

• Can be detected and measured with scanners


Radioisotopes (cont.)

• Radioisotopes are a valuable tool for biological

research and medicine

– Share same chemistry as their stable isotopes so

will be taken up by body

• Can then be used for diagnosis of disease

• All radioactivity can damage living tissue

– Some types can be used to destroy localized

cancers

– Some types cause cancer

• Radon from uranium decay causes lung cancer


2.3 Combining Matter

Molecules and Compounds

• Most atoms chemically combine with other

atoms to form molecules and compounds

– Molecule: general term for 2 or more atoms

bonded together

– Compound: specific molecule that has 2 or

more different kinds of atoms bonded together

• Example: C6H12O6

• Molecules with only one type of atom (H2 or O2) are

just called molecules


Mixtures

• Most matter exists as mixtures: two or more

components that are physically intermixed

• Three basic types of mixtures

– Solutions

– Colloids

– Suspensions


Figure 2.4 The three basic types of mixtures.


Solute

particles

Solute

particles

Solute

particles

Plasma

Settled

red blood

cellsUnsettled Settled

Example Example ExampleMineral water Jell-O Blood

Solute particles are very

tiny, do not settle out

or scatter light.

Solute particles are larger

than in a solution and

scatter light; do not

settle out.

Solute particles are very

large, settle out, and may

scatter light.

Solution Colloid Suspension

Mixtures (cont.)

• Solutions

– Are homogeneous mixtures, meaning particles

are evenly distributed throughout

– Solvent: substance present in greatest amount

• Usually a liquid, such as water

– Solute(s): substance dissolved in solvent

• Present in smaller amounts

• Example: blood sugar – glucose is solute, and blood

(plasma) is solvent


Mixtures (cont.)

• Solutions (cont.)

– True solutions are usually transparent

• Example: air (gas solution), salt solution, sugar

solution

• Most solutions in body are true solutions of gases,

liquids, or solids dissolved in water


Mixtures (cont.)

• Concentration of true solutions

– Three common ways to express concentrations:

1. Percent of solute in total solution

– How many parts of solute are in 100 total parts of

solution

– Solvent is usually water

– Example: 10 parts salt to 90 parts water is a 10% salt

solution

2. Milligrams per deciliter (mg/dl)

– Deciliter equals 1/100th of a liter

– Example: normal fasting blood glucose levels are

around 80 mg/dl


Mixtures (cont.)

3. Molarity (M) is number of moles of solute per liter of

solvent (water)

– 1 mole of a compound is equal to its molecular weight

(sum of atomic weights) in grams

– Example: glucose (C6H12O6 ) has a molecular wt of

180.12 amu, so 180.12 grams of glucose added to

enough H2O to make 1 liter is a 1 M solution of glucose

– 1 mole of any substance always contains 6.02 1023

molecules of that substance

– This number is called Avogadro’s number

– Molarities in the body are so small (can be 0.0001 M),

they are expressed in millimoles (mM) so

1000 mM = 1 M


Mixtures (cont.)

• Colloids

– Also known as emulsions; are heterogeneous

mixtures, meaning that particles are not evenly

distributed throughout mixture

• Can see large solute particles in solution, but these do

not settle out

• Gives solution a cloudy or milky look

– Some undergo sol-gel (solution to gel)

transformations

• Example: Jell-O goes from liquid to gel

• Cytosol of cell is also a sol-gel type solution


Mixtures (cont.)

• Suspensions

– Heterogeneous mixtures that contain large,

visible solutes that do settle out

– Example: mixture of water and sand

– Blood is considered a suspension because if left

in a tube, the blood cells will settle out


Difference Between Mixtures and Compounds

• Three main differences:

– Unlike compounds, mixtures do not involve

chemical bonding between components

– Mixtures can be separated by physical means,

such as straining or filtering; compounds can be

separated only by breaking their chemical bonds

– Mixtures can be heterogeneous or

homogeneous; compounds are only

homogeneous


2.4 Chemical Bonds

• Chemical bonds are “energy relationships”

between electrons of reacting atoms

– Chemical bonds are not actual physical

structures

• Electrons are the subatomic particles that are

involved in all chemical reactions

– They determine whether a chemical reaction will

take place and if so, what type of chemical bond

is formed


Role of Electrons in Chemical Bonding

• Electrons can occupy areas around nucleus

called electron shells

– Each shell contains electrons that have a certain

amount of kinetic and potential energy, so shells

are also referred to as energy levels

– Depending on its size, an atom can have up to 7

electron shells

– Shells can hold only a specific number of

electrons; the shell closest to nucleus is filled first

• Shell 1 can hold only 2 electrons

• Shell 2 holds a maximum of 8 electrons

• Shell 3 holds a maximum of 18 electrons© 2017 Pearson Education, Inc.

Role of Electrons in Chemical Bonding (cont.)

• Outermost electron shell is called valence shell

– Electrons in valence shell have the most

potential energy because they are farthest from

nucleus

– These are electrons that are involved in chemical

reactions


Role of Electrons in Chemical Bonding (cont.)

• Octet rule (rule of eights)

– Atoms desire 8 electrons in their valence shell

• Exceptions: smaller atoms (examples: H and He) want

only 2 electrons in shell 1

– Desire to have 8 electrons is driving force behind

chemical reactions

• Noble gases already have full 8 valence electrons (or

2 for He) so are not chemically reactive

– Most atoms do not have full valence shells

• Atoms will gain, lose, or share electrons (form bonds)

with other atoms to achieve stability of 8 electrons in

valence shell


Figure 2.5a Chemically inert and reactive elements.


Helium (He)

(2p+; 2n0; 2e−)

Neon (Ne)

(10p+; 10n0; 10e−)

2e 2e8e

Chemically inert elements

He Ne

Outermost energy level

(valence shell) complete

Figure 2.5b Chemically inert and reactive elements.


2e4e

2e8e

1e

1e

Na

2e6e

C

O

H

Chemically reactive elements

Hydrogen (H)

(1p+; 0n0; 1e−)

Oxygen (O)

(8p+; 8n0; 8e−)

Carbon (C)

(6p+; 6n0; 6e−)

Sodium (Na)

(11p+; 12n0; 11e−)

Outermost energy level

(valence shell) incomplete

Types of Chemical Bonds

• Three major types of chemical bonds

– Ionic bonds

– Covalent bonds

– Hydrogen bonds


Types of Chemical Bonds (cont.)

• Ionic bonds

– Ions are atoms that have gained or lost

electrons and become charged

• Number of protons does not equal number of

electrons



– Ionic bonds involve the transfer of valence shell

electrons from one atom to another, resulting in

ions

• One becomes an anion (negative charge)

– Atom that gained one or more electrons

• One becomes a cation (positive charge)

– Atom that lost one or more electrons

– Attraction of opposite charges results in an ionic

bond


Figure 2.6ab Formation of an ionic bond.


Sodium gains stability by losing one

electron, and chlorine becomes stable

by gaining one electron.

Sodium atom (Na)

(11p+; 12n0; 11e−)

Chlorine atom (Cl)

(17p+; 18n0; 17e−)

Sodium ion (Na+) Chloride ion (Cl−)

Sodium chloride (NaCl)

After electron transfer, the oppositely

charged ions formed attract each

other.

Na Cl Na Cl

+ −


• Most ionic compounds are salts

– When dry, salts form crystals instead of

individual molecules

– Example is NaCl (sodium chloride)


Figure 2.6c Formation of an ionic bond.


Large numbers of Na+ and Cl− ions

associate to form salt (NaCl) crystals.

Na+

Cl−


• Covalent bonds

– Covalent bonds are formed by sharing of two or

more valence shell electrons between two atoms

• Sharing of 2 electrons results in a single bond

• Sharing of 4 electrons is a double bond

• Sharing of 6 electrons is a triple bond

– Allows each atom to fill its valence shell at least

part of the time

– Two types of covalent bonds:

• Polar and nonpolar covalent bonds


Figure 2.7a Formation of covalent bonds.


Hydrogen atoms Carbon atom Molecule of methane gas (CH4)

or

H

H

H

H

H

H

H

H

C C

Resulting moleculesReacting atoms

Structural formula

shows single bonds.

Formation of four single covalent bonds: Carbon shares four electron pairs with four

hydrogen atoms.

Figure 2.7b Formation of covalent bonds.


or

Oxygen atom Oxygen atom Molecule of oxygen gas (O2)

Formation of a double covalent bond: Two oxygen atoms share two electron pairs.

O OOO

Structural formula

shows double bond.


Figure 2.7c Formation of covalent bonds.


or

Nitrogen atom Nitrogen atom Molecule of nitrogen gas (N2)

Formation of a triple covalent bond: Two nitrogen atoms share three electron pairs.

N NN N

Structural formula

shows triple bond.



• Covalent bonds (cont.)

– Nonpolar covalent bonds

• Equal sharing of electrons between atoms

• Results in electrically balanced, nonpolar molecules

such as CO2


Figure 2.8a Carbon dioxide and water molecules have different shapes, as illustrated by molecular models.


Carbon dioxide (CO2) molecules arelinear and symmetrical. They are nonpolar.


• Polar covalent bonds

– Unequal sharing of electrons between 2 atoms

– Results in electrically polar molecules

– Atoms have different electron-attracting abilities,

leading to unequal sharing

• Atoms with greater electron-attracting ability are

electronegative, and those with less are

electropositive



• Polar covalent bonds (cont.)

– H2O is a polar molecule

• Oxygen is more electronegative, so it exerts a greater

pull on shared electrons, giving it a partial negative

charge and giving H a partial positive charge

– Having two different charges is referred to as dipole


Figure 2.8b Carbon dioxide and water molecules have different shapes, as illustrated by molecular models.


V-shaped water (H2O) molecules have two poles of charge—a slightly more negative oxygen end (d−) and a slightly more positivehydrogen end (d+).

d+d+

d−

Figure 2.9 Ionic, polar covalent, and nonpolar covalent bonds compared along a continuum.


Ionic bond Polar covalent

bond

Nonpolar

covalent bond

Complete

transfer of

electrons

Separate ions

(charged

particles)

form

Unequal sharing

of electrons

Equal sharing of

electrons

Slight negative

charge (d−) at

one end of

molecule, slight

positive charge (d+)

at other end

Charge balanced

among atoms

Na+ Cl−

Sodium chloride Water Carbon dioxide


• Hydrogen bonds

– Attractive force between electropositive

hydrogen of one molecule and an

electronegative atom of another molecule

• Not true bond, more of a weak magnetic attraction

– Common between dipoles such as water

• What makes water liquid

– Also act as intramolecular bonds, holding a large

molecule in a three-dimensional shape


Animation – Hydrogen Bonds


Figure 2.10a Hydrogen bonding between polar water molecules.


O

O

O

O

H

H

H

H

H

H

H H

d−

d+

d+

d+

d+

d−

d−d−

d−

d+

d+

The slightly positive ends (d+) of the water

molecules become aligned with the slightly

negative ends (d−) of other water molecules.

Hydrogen bond

(indicated by

dotted line)

Figure 2.10b Hydrogen bonding between polar water molecules.


A water strider can walk on a pond because

of the high surface tension of water, a result

of the combined strength of its hydrogen

bonds.

2.5 Chemical Reactions

Chemical Equations

• Chemical reactions occur when chemical

bonds are formed, rearranged, or broken

• These reactions can be written in symbolic

forms called chemical equations

• Chemical equations contain:

– Reactants: substances entering into reaction

together

– Product(s): resulting chemical end products

– Amounts of reactants and products are shown in

balanced equations© 2017 Pearson Education, Inc.

Chemical Equations (cont.)

• Compounds are represented as molecular

formulas

– Example: H2O or C6H12O6

– Subscript indicates atoms joined by bonds

– Prefix denotes number of unjoined atoms or

molecules


Chemical Equations (cont.)

• Compounds are represented as molecular

formulas

– Example: H2O or C6H12O6 or H2 or CH4

– In chemical equations, subscripts indicate how

many atoms are joined by bonds, whereas prefix

means number of unjoined atoms (example: 4H)


Reactants

H + H →

4H + 1C →

Product

H2 (Hydrogen gas)

CH4 (Methane)

Types of Chemical Reactions

• Three main types of chemical reactions:

1. Synthesis (combination) reactions involve

atoms or molecules combining to form larger,

more complex molecule

• Used in anabolic (building) processes

A + B → AB


Figure 2.11a Types of chemical reactions.


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

Types of Chemical Reactions (cont.)

2. Decomposition reactions involve breakdown

of a molecule into smaller molecules or its

constituent atoms (reverse of synthesis

reactions)

• Involve catabolic (bond-breaking) reactions

AB → A + B


Figure 2.11b Types of chemical reactions.


Decomposition reactions

Bonds are broken in larger

molecules, resulting in smaller,

less complex molecules.

Example

Glycogen is broken down to release

glucose molecules.

Glycogen

Glucose

molecules


3. Exchange reactions, also called displacement

reactions, involve both synthesis and

decomposition

• Bonds are both made and broken

AB + C → AC + B

and

AB + CD → AD + CB


Figure 2.11c Types of chemical reactions.


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.

Adenosine triphosphate

(ATP)

Adenosine diphosphate

(ADP)

Glucose

Glucose-

phosphate

P P P

PP

P


• In living systems, these reactions are also

referred to as reduction-oxidation or redox

reactions

– Atoms are reduced when they gain electrons

and oxidized when they lose electrons

– Example: C6H12O6 + 6O2 → 6CO2 + 6H2O + ATP

• In this example, glucose is oxidized, and oxygen

molecule is reduced


Energy Flow in Chemical Reactions

• All chemical reactions are either exergonic or

endergonic

– Exergonic reactions result in a net release of

energy (give off energy)

• Products have less potential energy than reactants

• Catabolic and oxidative reactions

– Endergonic reactions result in a net absorption

of energy (use up energy)

• Products have more potential energy than reactants

• Anabolic reactions


Reversibility of Chemical Reactions

• All chemical reactions are theoretically

reversible

A + B ←→ AB

• Chemical equilibrium occurs if neither a forward

nor a reverse reaction is dominant

• Many biological reactions are not very reversible

– Energy requirements to go backward are too

high, or products have been removed


Rate of Chemical Reactions

• The speed of chemical reactions can be

affected by:

– Temperature: increased temperatures usually

increase rate of reaction

– Concentration of reactants: increased

concentrations usually increase rate

– Particle size: smaller particles usually increase

rate


Rate of Chemical Reactions

• Catalysts

– Catalysts increase the rate of reaction without

being chemically changed or becoming part of

the product

– Enzymes are biological catalysts


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Chemistry Comes Alive - Dr. Jerry Cronindrjerrycronin.weebly.com/uploads/5/9/7/4/5974564/... · 2.3 Combining Matter Molecules and Compounds •Most atoms chemically combine with