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Matter and Energy

• Matter—anything that occupies space and has

mass (weight)

• Energy—the ability to do work

– Chemical

– Electrical

– Mechanical

– Radiant

Composition of Matter

• Elements—fundamental units of matter

– 96% of the body is made from four elements

• Carbon (C)

• Oxygen (O)

• Hydrogen (H)

• Nitrogen (N)

• Atoms—building blocks of elements

Identifying Elements

• Atoms are composed of 3 particles

– Protons – mass 1, charge +

– Neutrons – mass 1, no charge

– Electrons – mass 0, charge -

• Atomic number—equal to the number of protons that the atom contains.

This defines which element it is. For example, any atom with 6 protons is

Carbon. Any atom with 4 protons is Beryllium.

• Atomic mass number—sum of the protons and neutrons. This is how

“heavy” the atom is. Only protons and neutrons have significant “weight”

electrons are almost massless.

Atomic Structure• Nucleus

– Protons (p+)

– Neutrons (n0)

• Outside of

nucleus

– Electrons (e-)

Figure 2.1

Atomic Structure of Smallest Atoms

Figure 2.2

Isotopes and Atomic Weight• Isotopes

– Have the same number of protons

– Vary in number of neutrons

Figure 2.3

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Isotopes and Atomic Weight

• Atomic weight

– Close to mass number of most abundant isotope

– Atomic weight reflects natural isotope variation

2.1 What Are the Chemical Elements

That Make Up Living Organisms?

Radioisotopes are unstable, they give off

energy in the form of alpha, beta, and

gamma radiation from the nucleus.

This radioactive decay transforms the atom,

including changes in the number of

protons.

The halflife describes the amount of time it

takes an element to decay

Radioisotope Half-life

Phosphorus - 32

Tritium

Carbon - 14

Potassium - 40

Uranium – 238

14.3 days

12.3 years

5,700 years

1.3 billion years

4.5 billion years

2.1 What Are the Chemical Elements

That Make Up Living Organisms?

Energy from radioactive decay can interact

with surrounding material.

Radioisotopes can be incorporated into

molecules and act as a “tag” or label.

Electrons and Bonding

• Atoms are united by chemical bonds…

• Electrons occupy energy levels called electron

shells

• Electrons closest to the nucleus are most strongly

attracted

• Each shell has distinct properties

– The number of electrons has an upper limit

– Shells closest to the nucleus fill first

Electrons and Bonding

• Bonding involves interactions between

electrons in the outer shell (valence shell)

• Full valence shells do not form bonds

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Inert Elements

• Atoms are stable (inert) when the

outermost shell is complete

• How to fill the atom’s shells

– Shell 1 can hold a maximum of 2 electrons

– Shell 2 can hold a maximum of 8 electrons

– Shell 3 can hold a maximum of 18 electrons

Inert Elements

• Atoms will gain, lose, or share electrons to

complete their outermost orbitals and

reach a stable state

• Rule of eights

– Atoms are considered stable when their

outermost orbital has 8 electrons

– The exception to this rule of eights is Shell 1,

which can only hold 2 electrons

Inert Elements

Figure 2.5a Figure 2.5b

Reactive Elements

• Valence shells are not full and are unstable

• Tend to gain, lose, or share electrons

– Allow for bond formation, which produces stable

valence

Chemical Bonds

• Ionic bonds

– Form when electrons are completely transferred

from one atom to another

• Ions

– Charged particles

• Anions are negative

• Cations are positive

• Either donate or accept electrons

Ionic Bonds

Figure 2.6

+ –

Sodium atom (Na)(11p+; 12n0; 11e–)

Chlorine atom (Cl)(17p+; 18n0; 17e–)

Sodium ion (Na+) Chloride ion (Cl–)

Sodium chloride (NaCl)

ClNaClNa

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Chemical Bonds

• Covalent bonds

– Atoms become stable through shared electrons

– Single covalent bonds share one pair of electrons

– Double covalent bonds share two pairs of

electrons

Examples of Covalent Bonds

Figure 2.7b

Examples of Covalent Bonds

Figure 2.7c

Polarity• Covalently bonded

molecules

– Some are non-polar

• Electrically neutral

as a molecule

– Some are polar

• Have a positive and

negative side

Figure 2.8

How Do Atoms Bond to Form

Molecules?

Electronegativity: the attractive force that

an atomic nucleus exerts on electrons

Electronegativity depends on the number of

positive charges (protons) and the

distance between the nucleus and

electrons.

Table 2.3

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How Do Atoms Bond to Form

Molecules?

Polar molecules that form hydrogen bonds

with water are hydrophylic (“water-

loving”).

Nonpolar molecules such as hydrocarbons

that interact with each other, but not with

water are hydrophobic (“water-hating”).

Chemical Bonds

• Hydrogen bonds

– Weak chemical bonds

– Hydrogen is attracted to the negative portion of

polar molecule

– Provides attraction between molecules

Hydrogen Bonds

Figure 2.9

Table 2.1

2.2 How Do Atoms Bond to Form

Molecules?

van der Waals forces: attractions between

nonpolar molecules. They result from

random variations in electron distribution.

Individual interactions are brief and weak,

but summed over a large molecule, can be

substantial.

Patterns of Chemical Reactions

• Synthesis reaction (A + B�AB)

– Atoms or molecules combine

– Energy is absorbed for bond formation

– Ex: dehydration synthesis/condensation reaction

• Decomposition reaction (AB�A + B)

– Molecule is broken down

– Chemical energy is released

– Ex: hydrolysis

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Synthesis and Decomposition

Reactions

Figure 2.10a

Synthesis and Decomposition Reactions

Figure 2.10b

pH Scale, Measurements and Interactions

pH Scale: 0-14

acids: 0-6

neutral: 7

base: 8-14

pH Scale, Measurements and Interactions

When acids dissolve in water, they release hydrogen ions—H+ .

H+ ions can attach to other molecules and change their properties.

High concentration of H+ = acidic

(Bases accept H+ ions.)

−+

+→ ClHHCl

2.4 What Properties of Water Make It

So Important in Biology?

Organic acids have a carboxyl group:

Weak acids: not all the acid molecules

dissociate into ions.

+−

+−→− HCOOCOOH

When bases dissolve in water, they release hydroxide ions—OH-.

OH- ions can attach to other molecules and change their properties AND accept H+.

High concentration of OH- = basic

Any compound that accepts H+ is a base (strong or weak)

2.4 What Properties of Water Make It

So Important in Biology?

−+

+→ OHNaNaOH

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2.4 What Properties of Water Make It

So Important in Biology?Weak bases:

• Bicarbonate ion

• Ammonia

• Compounds with amino groups

323COHHHCO →+

+−

++

→+43

NHHNH

++

→+−32

NHHNH

2.4 What Properties of Water Make It

So Important in Biology?Water is a weak acid.

−+

+→ OHHOH2

2.4 What Properties of Water Make It

So Important in Biology?pH = negative log of the molar concentration

of H+ ions.

H+ concentration of pure water is 10–7 M, its pH = 7.

Lower pH numbers mean higher H+

concentration, or greater acidity.

2.4 What Properties of Water Make It

So Important in Biology?Living organisms maintain constant internal

conditions, including pH.

Buffers help maintain constant pH.

A buffer is a weak acid and its corresponding

base.

323COHHHCO →+

+

Buffer examples

• In the blood these compounds buffer pH

• Carbonic acid: H2CO3

• Bicarbonate ions: HCO3-

• HCO3- + H+ � H2CO3

Figure 2.17 Buffers Minimize Changes in pH