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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
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