Basic Chemistry

Chemical Elements

• basic unit of matter – 92 recognized elements – 25 essential for life – 6 major elements in living

organisms • carbon

• hydrogen

• nitrogen

• oxygen

• phosphorus

• sulfur

CHNOPS ( 98% )

atom - smallest unit of matter that retains the properties of an element - carbon is an element - may have multiple atoms of

carbon (only one element)

Subatomic particles:

Particle Location Charge Mass

Proton nucleus positive1 AMU

Neutron nucleus neutral1 AMU

Electron orbiting negative negligible nucleus

126

Carbon

CMassMassNumberNumber

AtomicAtomicNumberNumber

AtomicAtomicSymbolSymbol

Atomic Symbols and the Periodic Table

atomic # is the # of protons (unique to each element)

Atoms begin electrically neutral thus:positive and negative charges balance out # of protons = # of electrons

Mass # equals the # of protons + # of neutrons

Each element has a unique # of electrons# of electrons --> influences interactions w/ other atoms

Isotopes:

atoms typically have equal # of protons and neutrons - except hydrogen which only has one proton

atoms MAY vary in the # of neutrons they have (Isotopes) - if not equal with # of protons = more unstable - radioactive isotopes

- acts the same as stable isotopes but can be ‘traced’

- used in medical diagnosis and research

- spontaneously break down at a constant rate

- can be used in dating fossils

Vertical columns indicatenumber of electronsin outermost shell

1

H1.008

3

Li6.941

11

Na22.99

19

K39.10

4

Be9.012

12

Mg24.31

20

Ca40.08

5

B10.81

13

Al26.98

21

Ga69.72

6

C12.01

14

Si28.09

22

Ge72.59

7

N14.01

15

P30.97

23

As74.92

8

O16.00

16

S32.07

24

Se78.96

9

F19.00

17

Cl35.45

25

Br79.90

10

Ne20.18

18

Ar39.95

26

Kr83.60

2

He4.003

1

2

3

4

Horizontal periods indicate

total number

of electron shells

I

II III IV V VI VII

VIII

Periodic

Table

Electrons have energy: - electrons are attracted to positively charged nucleus - energy absorbed pushes the electron away

creates potential energy - energy is released when electron moves closer

releases energy for use in chemical reactions

The Octet Rule forDistribution of Electrons

• Bohr models show electron shells as concentric circles around nucleus

– Each shell has two or more electron orbitals

• Innermost shell has two orbitals

• Others have 8 or multiples thereof

• Atoms with fewer than 8 electrons in outermost shell are chemically reactive

– If 3 or less – Tendency to donate electrons

– If 5 or more – Tendency to receive electrons

Bohr Models of Atoms

Atoms interact and form chemical bonds dependent upon electron configuration.

- electrons arranged in shells - # empty spaces in outer shell = the # of potential binding sites

open circle - available binding sites

Molecules:

• are formed when two or more atoms are bound together • made up of different elements or the same element

• If all atoms in molecule are of the same element Material is still an element 2 oxygen atoms form 1 molecule oxygen gas

•If at least one atom is from a different elementMaterial formed is a compound

2 hydrogen & 1 oxygen form 1 water moleculeCharacteristics dramatically

different from constituent elements

Molecular / Chemical formula used to represent the # of atoms of each type of element in a molecule:

CH4 - means one carbon and four hydrogen atoms together

H2O - means two hydrogen and one oxygen

Chemical Bonds: chemical bonds hold atoms together in a molecule & influence how molecules interact with one another

Types of Bonds between atoms/elements:

1. Ionic bonds transfer of electrons between two atoms

2. Covalent bonds sharing of electrons between two atoms

- equal sharing - non-polar covalent bonds

- unequal sharing - polar covalent bonds

3. Hydrogen bondsweak electrostatic charge - opposites attract

bond between polar molecules or polar areas of a large molecule

Ionic Bonds:

• reactivity - tendency to lose or gain electrons

• full outer valence shell = stability

IONS: Na+ / Cl-

electrically charged atoms

• transferring electron

• formation of IONS

IONIC BOND:

• ions -- opposites attract

• ionic bond - moderate strength bond

• Na+ & Cl- form NaCl (sodium chloride)

• fully charged atoms (lost or gained an electron)

• strong (and opposite) electrical charges attract one another

fig 2.7

Covalent Bonds:

• sharing electrons

• # of binding sites

• dependent on outer electron shell

• atoms share ONE pair of electrons

• single covalent bond

• atoms share TWO pairs of electrons

• double covalent bonds

double covalent bonds harder to break

fig 2.8

Non-polar Covalent Bonds:

• equal sharing of electron

• equal pull on electron

• molecule electrically neutral

• covalent bonds share electrons

• unequal sharing

• electron spends more time around one of the atoms in bond

• molecule has slightly negative side & slightly positive side

examples: H2, CH4

examples: H2O, NH3

Polar covalent bonds:

fig 2.8

fig 2.9

Hydrogen bonds:

• electronegative atom

(an atom with a slightly stronger pull on a shared electron)

attracts a HYDROGEN atom already w/in a polar covalent bond in a different molecule or another part of the same molecule

polar molecule

excellent example: water

fig 2.9b

Ionic vs. Hydrogen bonds:

Ionic: ions (charged atoms) attracted to one another - strong electrical charge

Hydrogen: two polar molecules or parts of molecules attracted to one another - slight electrical charge

stronger bond than hydrogen

Properties of Water:

Water is a polar molecule due to polar-covalent bonds that are created.

As such, water is held together by hydrogen bonds.

This polarity and subsequent hydrogen bonding creates a number unique properties that are essential to life on Earth.

1. Its heat capacity2. Its cohesive nature3. Its reaction when forming a solid (ice)4. Its use as a solvent5. Its ability to ionize

1. Temperature Buffer -- Heat Storage

water changes temperature slowly

heat energy is absorbed to break bonds

bonds formed before movement of molecule slows

as water is heated…. heat breaks the bonds between water molecules before water can vaporize!

as water is cooled…. bonds are formed first heat released as converted to solid Temperature

regulation in humans: •evaporative coolingevaporative cooling

Water has a high heat capacity- Temperature = rate of vibration of molecules

- Apply heat to liquid- Molecules bounce faster- Increases temperature

- But, when heat applied to water - Hydrogen bonds restrain bouncing

- Temperature rises more slowly per unit heat

- Water at a given temp. has more heat than most liquids

Thermal inertia – resistance to temperature change- More heat required to raise water one degree than

most other liquids (1 calorie per gram)- Also, more heat is extracted/released when lowering

water one degree than most other liquids

High heat of vaporization

- To raise water from 98 to 99 ºC; ~1 calorie- To raise water from 99 to 100 ºC; ~1 calorie- However, large numbers of hydrogen bonds

must be broken to evaporate water- To raise water from 100 to 101 ºC; ~540

calories!

This is why sweating (and panting) cools

- Evaporative cooling is best when humidity is low because evaporation occurs rapidly

- Evaporative cooling works poorest when humidity is high because evaporation occurs slowly

Heat of fusion (melting)

- To raise ice from -2 to -1 ºC; ~1 calorie- To raise water from -1 to 0 ºC; ~1 calorie- To raise water from 0 to 1 ºC; ~80 calories!

This is why ice at 0 ºC keeps stuff cold MUCH longer than water at 1 ºC

This is why ice is used for cooling

- NOT because ice is cold- But because it absorbs so much heat before it will warm by one degree

Heat Content of Waterat Various Temperatures

2. Cohesion and Adhesion of Water

- Cohesion – Hydrogen bonds hold water molecules tightly together

- Adhesion – Hydrogen bonds form between water and other polar materials

- Hydrogen bonds -- breaking & reforming - bonds last a few trillioniths of a second!

- always substantial % bonded to neighboring molecules

High Surface Tension- Water molecules at surface hold more tightly than below surface- Amounts to an invisible “skin” on water surface- Allows small nonpolar objects (like water-

strider) to sit on top of water

Allows water be drawn many meters up a tree in a tubular vessel

3. Ice Formation- water molecules densest at 4°C

- as temp drops bonds become more spaced & stable - density drops- frozen water less dense than liquid water

Density of Waterat Various Temperatures

A Pond in WinterLakes/oceans don’t stay frozen over because ICE floats

Ice acts as an insulator on top of a frozen body of water

Otherwise, oceans and deep lakes would fill with ice from the bottom up

Melting ice draws heat from the environment

Turnover of nutrients / oxygen when lake thaws

4. Water is the universal solvent

- A solvent (the most abundant part) and- A solute (less abundant part) that is dissolved

in the solvent

- Polar compounds readily dissolve; hydrophilic- Nonpolar compounds dissolve only slightly; hydrophobic

- Ionic compounds dissociate in water- Na+

Attracted to negative (O) end of H2OEach Na+ completely surrounded by

H2O

- Cl-Attracted to positive (H2) end of H2OEach Cl- completely surrounded by

H2O

Hydrophilic Hydrophilic molecules: molecules: water-lovingwater-loving

break-up / surround polar molecules and ionized compounds

pH - measure of hydrogen ions (H+)

concentration:

• most chemical reactions w/in our bodies influenced by pH

• most biological fluids act as buffers - neutralizing pH

• basic solutions (high pH) - lower H+ concentration

• acidic solutions (low pH)- high H+ concentration

5. Water Ionizes

water molecules sometimes break

• H2O <--> OH- + H+

AcidsDissociate in water and release hydrogen ions (H+)Sour to tasteHydrochloric acid (stomach acid) is a gas with symbol HCl

In water, it dissociates into H+ and Cl-Dissociation of HCl is almost total, therefore it is a strong acid

Bases:Either take up hydrogen ions (H+) or release hydroxide ions (OH-)Bitter to tasteSodium hydroxide (drain cleaner) is a solid with symbol NaOH

In water, it dissociates into Na+ and OH-Dissociation of NaOH is almost total, therefore it is a strong base

pH scale used to indicate acidity and alkalinity of a solution.Values range from 0-14

0 to <7 = Acidic7 = Neutral>7 to 14 = Basic (or alkaline)

The pH Scale

Health of organisms requires maintaining pH of body fluids within narrow limits

- Human blood normally 7.4 (slightly alkaline)

- Many foods and metabolic processes add or subtract H+ or OH- ions

- Reducing blood pH to 7.0 results in acidosis

- Increasing blood pH to 7.8 results in alkalosis

- Both life threatening situations

- Bicarbonate ion (-HCO3) in blood buffers pH to 7.4