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Phys Sci Week 6 Basic Chemistry
and Electricity
• Fundamental Forces involved
– Strong and Weak Nuclear Forces
• Keeps nucleus together
• Electromagnetic
– All chemical Reactions and Electricity involves
the EM force
• It is the electron interaction that actually causes
chemical reactions and electricity.
Models
• (A model is some working representation of a
real system
• Models are simplifications or idealizations of the
real thing that makes a theory useful.
• Some types of models:
–
–
–
–
–
–
–
Parables
Rules
Diagrams - schematics
Physical models "like a world globe"
Analogs
Mathematical
Computer
Atomic Models
• All atomic models are mathematical
models that can be represented by a
diagram.
• Pre-atomic: mater matter was thought
to be continuous
• The following slides will give examples
of atomic models
Definitions
Module 1 Study Guide Questions p 24 # 1 -14
• Atom - atomic - smallest particle to make up an
element
– Elements all on type of atom
• Sub-atomic - parts of atoms
– Electrons, Protons, Neutrons, Quarks & Strings
• Molecule - smallest particle of a compound
– Composed of two or more atoms bonded together by
sharing or an exchange of electrons.
– Atoms that make up a molecule can be of the same type
or different types.
• Ion - charged atom or molecule
• Isotope - same number of proton - same atom type different number of neutrons
History of the study Chemistry &
Electricity
• Both start with the idea of the
atom in 460 to 370 B.C, meaning
“composed of particles which
cannot be divided."
• Until early 20th century people
believed that the smallest form of
matter was an atom,
– imagined to be the shape of a ball
(a sphere).
– Now, we know that atoms are
made of even smaller parts.
• Modern idea of the atoms was
formulated in the early 1900s.
Rutherford's' Model
1920
Dalton's Atomic Model
1800
1900
Picture of Dalton's Theory of the atom
Thomson's Model (1890 - 1920)
• Discovered of the electron experimenting with cathode (negative
electrode) rays - electrons added to the atomic model.
Rutherford Model (1880's - 1930)
Discovered the proton
•
Diagram of Rutherford's Model
Current Bohr Model ~1939
The Structure of an Atom
• Each atom is
made up of:
– Protons and
Neutrons in the
nucleus
– Electrons on
the outside of
the atom
Bohr Model
Energy Levels - Orbits
• Electron's can only exist at
certain energy levels or
quanta
• Photons (packet of light
energy) given off when
electrons change energy
state
• Beginning of quantum
theory which explains very
small - remember theory of
gravity can not explain the
very small
Current Bohr Model
Chadwick Discovery - The Neutron:
The Proton - Neutron Model
•
Chadwick discovers the neutron as part of the nuclease in 1932. Led
to current model which is still in use.
More on Bohr Model on p 314
• Relative size of atom and parts - p 315
• If the electron orbit were the size of as baseball
stadium - nucleus would be the size of a marble
• Atom is 99.99 % empty space (p 316)
– When we "touch" something we really are feeling the
Interaction of the negative repulsive forces of the
electrons (p316)
A closer look at quarks
• Protons and neutrons made up of three
quarks where the exchange of particle
are called gluons
The Strong Force (p 325)
• Very strong short range force that keeps the
nucleus together
– P to P, N to N and P to N.
– Note that as atoms get heavier the number of
neutrons exceeds the number of protons - this
provides more strong force to keep the Protons from
moving apart due to the repulsive + charges
repelling.
• Exchange of the pion particle causes the strong
force
– Pion is a short lived particles - cause very strong
forces - p 326
Details of Parts of the Atom
• Nucleus: Made up of Protons and Neutrons
– Weigh about the same both (about 1 amu)
– Neutron slightly heavier
– Proton + charge neutron - charge
• Electron: about 2000 times lighter than Proton or
Neutron - Negative charge
• Ions occur when there is an unbalanced charge
due to a lack of (+ ion) or abundance of (- ion)
• Quarks make up Protons, and Neutrons
– Breakup of neutron into proton, electron and
antineutrino
Atomic Number and Mass
• Atomic number - number of protons dictates type of atom (p317)
• Mass number (AMU) = Sum of Protons and
Neutrons (see figure 13.3 on page 318)
• Isotopes: Same number of protons - but
different number of neutrons - same type of
atom with a different atomic mass (mass
number) (Page 18-319)
Current Quantum Mechanical
Models
• Based upon
– No two particles can occupy the same space (have the
same quantum numbers)
– The uncertainty principal (can't no both positions and
velocity of a particle at the same time)
• Bohr Orbits versus orbital, electron levels, or
shells
• Orbit number and Electron capacity
– 1 = 2, 2 = 8, 3 = 18, 4 = 32, 5 = 50
The Periodic Table (p 320)
• Elements (p 319 - A collection of atoms
that have all the same number of
protons - made up of the same type of
atoms.
• Modern Periodic Table (p 103)
– Mendelev's 1880's - based upon atomic
mass - Dalton's Atomic Model
– Mosely 1912 - based upon atomic number
Select an element
(
= Internet link )
Parts of Periodic Table (p 321)
• Groups or Family - arranged in columns
– Have similar properties because they have same
number of valence electrons
– Similar electron configuration (indicated by
Roman numeral).
• Main group A
• Transition metals B group
• Periods or series are in rows
–
–
–
–
Metals
Metalloids
Nonmetals
Lanthanide Series and Actinide Series (p107)
Periodic Trends
• Predicating Electron Configurations
• Atomic and ionic radii
–
–
–
–
Decrease in size as you move from left to right (gets heavier)
Increase in size as you move down a column
Negative ions increase in size
Positive ions decrease in size.
• Ionization Energy: Energy needed to remove electron
– Increase left to right
Decrease for the heavier atoms (down columns)
• Size constant for metals
Use the periodic table to answer the
following questions about Iron (Fe):
• What is the atomic number of Iron?
• How many protons and electrons does a Fe
atom have?
• About how many neutrons would a Fe atom
have?
• If a Fe atom were to lose one electron, it would
have an electrical charge of
Use the periodic table to answer the
following questions about Iron (Fe):
• What is the atomic number of Iron? 26
• How many protons and electrons does a Fe
atom have?
• About how many neutrons would a Fe atom
have?
• If a Fe atom were to lose one electron, it would
have an electrical charge of
Use the periodic table to answer the
following questions about Iron (Fe):
• What is the atomic number of Iron? 26
• How many protons and electrons does a Fe
atom have? 26
• About how many neutrons would a Fe atom
have?
• If a Fe atom were to lose one electron, it would
have an electrical charge of
Use the periodic table to answer the
following questions about Iron (Fe):
• What is the atomic number of Iron? 26
• How many protons and electrons does a Fe
atom have? 26
• About how many neutrons would a Fe atom
have? 26
• If a Fe atom were to lose one electron, it would
have an electrical charge of
Use the periodic table to answer the
following questions about Iron (Fe):
• What is the atomic number of Iron? 26
• How many protons and electrons does a Fe
atom have? 26
• About how many neutrons would a Fe atom
have? 26
• If a Fe atom were to lose one electron, it would
have an electrical charge of -1
Radioactivity - the Weak Force (p 327)
• Radioactivity or radioactive decay Breakdown
of an atom into two or more atoms plus energy
• Particles are given off
caused by a release of energy when the weak
nuclear force is released – Weak force is similar to the EM force (Weak force is
now thought to be a different form of the EM force analogous of how water can be ice, liquid or vapor)
•
•
Radioactive
isotopes
Isotope of an atom that is radioactive - used in medicine.
Types of radioactive decay (Radioactivity)
–
–
•
Alpha decay - 2N + 2 P (He nucleus leaves the atom) + energy
–
–
•
Nucleus gives off high energy called gamma rays
Example (p 329) Th 239 Unstable -> Th 239 stable + gamma ray (photon) both have 90 P and 139 N.
Dangers of Radioactivity (p 330)
–
–
–
–
•
Example Po 214 (84P) -> Pb 210 (82P) + alpha particle (He nucleus) +E
See Fig13.56 page 328
Gamma decay
–
–
•
Beta decay - Neutron -> proton + electron (beta particle) energy
Example U239(92P) -> NP239(93P)
Like tiny bullets - penetrates below the skin
Gamma light but fast - most damaging - takes a lead shield to stop them
Beta light - faster than alpha but slower than gamma least damaging. - Thin metal stops them
Alpha slow but heavy = paper stop them
Rate of Radioactive Decay (p 332)
–
Some radioactive elements undergo radioactive decay quickly, some very slowly
–
Half life is the time it takes for half of radio active material to decay - example 10 gram of U239(92P) -> 5 gram of
U239(92P) = 5 grams of NP239(93P (see fig 13.6 page 333)
Radioactive Dating (p334)
Using the amount of radioactive material in a substance to detriment age based upon decay rates
• Example C-14
C14 decay to C12
Half life is 5700 years
Assumption is that when organisms died it had a certain amount of C14 in it so there fore we can tell when
it died because it would stop taking in C14 and it would have so much less C14 so if it stated out with 10
grams of C14 and it now has 1 gram of C14 then it would be 50,000 years
Problems with C14 - deductive part - assumption of how much C14 was in the organism to start no on really
knows. Based upon uniformatariansim - about value since 1945 have change - why?? What does this tell
us.
–
Chemical Bonds:
• Sharing or “borrowing” outer shell – valence – electrons.
– Follow rule of the octave
• S - , P 8, D 8 and so on
– note –electron with proton is intra-molecular interactions
– Intermolecular interaction - Example Na+ Cl-
• Ionic bonds – borrowing electrons – not really consider a bond, but an
ionic attraction'
• Covalent Bonds - sharing of electrons – true bond – very strong bonds
–
–
–
–
Intermolecular Covalent bonds
Single Bond
Double Bond
Triple Bond
• Vader Walls – Hydrogen Bonds – weak interactions – not a true bonds
cases by
– permanent dipole–permanent dipole forces
– permanent dipole–induced dipole forces
– induced dipole-induced dipole
Ionic Bond
Molecules and Chemical
Compounds (AP p 134 – 136)
• Single atoms Monatomic: In physics and chemistry, monatomic is a
combination of the words
– "mono" and "atomic," and means "single atom." It is usually applied to
gases: a monatomic gas is one in which atoms are not bound to each other.
– At standard temperature and pressure (STP), all of the noble gases are
monatomic. These are helium, neon, argon, krypton, xenon and radon. The
heavier noble gases can form compounds, but the lighter ones are
unreactive.
– All elements will be monatomic in the gas phase at sufficiently high
temperatures.
• Molecules: Molecules are formed when atoms linked together (AP 134 –
135)
• Diatomic molecules are molecules composed only of two atoms, of
either the same or different chemical elements. The prefix di- means two
in Greek. Common diatomic molecules are hydrogen, nitrogen, oxygen,
and carbon monoxide. Most elements aside from the noble gases form
diatomic molecules when heated, but high temperatures - sometimes
thousands of degrees - are often required.
Forming Covalent Molecules Octet Rule
• Octet rule is a chemical rule of thumb that
states that atoms tend to combine in such
a way that they each have eight electrons
in their valence shells, giving them the
same electronic configuration as a noble
gas.
Chemical Formulas (p 85)
• Chemical Symbols and Formulas of Compounds
– Use of subscript - goes with prior symbol
– Use of coefficient - in front of atom or compound·
• Using symbols to represent chemicals such as
– H 2O
CO2
CH4
2H2O
• Using symbols to represent chemical Reactions
such as
• 2H2 + O2
2H2O
• Where names come from - some Latin ferrum so Fe
for Iron.
Two basic types of molecules
a. Inorganic molecules are molecules that do
not contain carbon. Inorganic molecules
make up things such as rocks, minerals and
metals. Many gases such oxygen, nitrogen
and hydrogen are composed of inorganic
molecules. Water is also an inorganic
molecule.
• b. Organic molecules are molecules that do
contain carbon. Organic molecules make up
living things, hydrocarbons such as coal and
oil, and liquids such things alcohol.
The C12 Bohr Model -
Division of Matter
Matter
Mixtures
Heterogeneous Homogeneous
Pure Substances
Elements Compounds
Compound and Mixtures
• a. Compound: Two or more elements that are
chemically combined and can not be
separated by physical methods. For example
water is a compound.
• b. Mixture: Two or more elements that are
blended together but could be separated by
physical means. For example water is a
compound.
• c. Key difference between them: Mixture can be
and compounds cannot be separated by
physical means.
Name Compound (C) or Mixture (M)
•
•
•
•
a. Water ______(C)_____________
b. Vinegar _____(M)______________
c. Sand _______(M)____________
d. Salt ______(C)_____________
Types of Mixtures
• Two basic groups:
– Heterogeneous mixtures are not spread out evenly. Example: a
bottle of liquid salad dressing, where the water and oil separate.
– Homogeneous mixtures substances are spread evenly
throughout. A homogeneous mixture is called a solution.
Example: vinegar (water and acetic acid are mixed evenly
throughout). Other examples: sea water, soft drinks, and glass
• Classes of Mixtures
– Solutions is a homogeneous mixture in which one substance (the
solute) is dissolved in another substance (the solvent). Example: salt
water (Water, the solvent, plus salt, the solute, produces the solution of
salty water.)
– Suspensions - a heterogeneous mixture in which the particles are
large enough to be seen by a microscope or the unaided eye
(eventually, they settle out of the mixture). Example: stirring a teaspoon
of dirt in a glass of water.
– Colloids - a mixture where the sizes of particles in the mixture are
between those of a solution and a suspension. The particles in a colloid
appear evenly distributed. Examples: fog, cheese, butter, jellies,
whipped cream.
• . Label the following mixtures as a solution
(SOL), suspension (SUP), or colloid (COL).
• Name of Mixture solution (SOL), suspension
(SUP),
• or colloid (COL)
• a. Fog _(COL)_________
• b. Salt mixed in water
_________(SOL)__________
• c. Sand mixed in water
____(SUP)_______________
Chemical and Physical Properties
and changes (AP p 136 – 137)
•
Physical properties can be observed or measured without changing the composition of
matter. Physical properties are used to observe and describe matter. Physical properties
include:
–
appearance, texture, color, odor, melting point, boiling point, density, solubility, polarity, and many
others.
•
Chemical properties of matter describes its "potential" to undergo some chemical change
or reaction by virtue of its composition. What elements, electrons, and bonding are
present to give the potential for chemical change. It is quite difficult to define a chemical
property without using the word "change". Eventually you should be able to look at the
formula of a compound and state some chemical property.
•
•
Chemical and Physical Changes
Physical changes occur when objects undergo a change that does not change their
chemical nature. A physical change involves a change in physical properties. Physical
properties can be observed without changing the type of matter. Physical changes are
reversible.
•
Examples of physical properties include: texture, shape, size, color, odor, volume, mass,
weight,
and density.
•
•
Chemical changes are the changes in a substance through chemical reactions. The
chemical reactants form a new product with equal mass.
Wonder of Water (p 81)
• 4. Look at its composition H2O - what would you think it would
be based upon its
• molecular weight?
• 5. Why:
• 6. Is it liquid at normal temperatures - needed for life
• 7. What can it change phase - weather - keeping balance temp
• 8. Why does it have such a high heat capacity for a simple
molecule
• 9. why does it expand when it freezes
• 10. Why can it hold more O2 when it gets colder 2
• 11. Evidence of the God as a Creator of the universe and His
Love
The Composition of Water (p 81)
• · Water is made up 2 H for every O atom H2O
• · Discover through the process of Electrolysis
• · Pass current through a substance (water)
breaks substance down
• · Negative tem H gas (H slight positive) Negative
terminal O2 Gas
• · Water give of H2 and O2 gas in a 2:1 ratio
• Experiment 4.1 The chemical composition of
water (page 81)
Water's Polarity (p 86)
•
•
•
•
•
•
•
•
•
•
•
•
Look at figures 4.2 and 4.3
H end of the molecule is slightly
positive
O end of molecule slightly negative
· Polar Molecule: Water has polarity (+
and - ends) and is called a polar
molecule
- most molecules have some polar
qualities. Water has just enough to
give its
special properties.
· Non Polar Molecules: Some
molecules are quite nonpolar like oil
which don't
mix well with water.
· What's the big deal it is just water - if
you gave someone a great gift and
they
scoffed at how would you feel. Some
substance allow water an oil to be
soluble in
both soap.
· Experiment 4.2 Waters polarity (p 86)
(also water oil and dish detergent.
Water as a Solvent (p 90)
•
•
•
•
•
•
•
Solution: when you dissolve a solid or liquid into a liquid to form
Solvent - A liquid substance capable of dissolving other substances.
Water called near universal solvent
Solute - A substance that is dissolved in a solvent solid or liquid
Ionic compounds (e.g NaCl) - water dissolves well because they are polar
molecules. (see figure 4.5.
Experiment 4.3: Solvents and Solutes (p 90)
Hydrogen Bonding (p93)
• ·Weak bond of hydrogen on one
molecule with Oxygen of another
molecule.· See Fig 4.6. (p 94)
• Hydrogen bonds link molecules together
(related to polar nature of H2).
– See special statement on water bottom
of page 94
• Gives water its special properties
– Latent heat - Phase change
– High heat capacity
– Liquid when you would think it would be
a gas
– Why it forms a crystalline structure and
explains when it freezes
• ·Cohesiveness of water
• Exp: Comparing solid water to solid
butter (p 95)
Water's Cohesion (p 97)
•
•
•
•
· The tendency of water to stick together
· Causes surface tension
· Meniscus shape of water on a glass - in nature xylem
· What it is hard to get all the water of something as compared to
alcohol which is
• more nonpolar.
• Exp Water Cohesion (p 97)
• Exp The forces between Molecules (p99)
Hard and Soft Water (p100)
• · Hardware has dissolved ions of Ca+ or Mg+ in it.
• · Does not soap up as easily
• · Can soften water by replacing Ca+ with Na+ - but it is not as
healthy.
End
• Experimental Terminology
• · Experimental Error - Errors/mistakes
cause value to not be perfect
• · Peer review - other scientist look at
results in Journal
• · Example cold versus hot fusion.
Greek: (Democritus 440 BC)
• The Greeks where the first to have the
idea that matter is made up of discreet
fundamental particles that can't be
divided.
• Atoms can only combine in certain
ratios - Law of definite composition.
H2O, H2O2
Dalton: First Experimental
Model 1770 - 1840. (see page 69)
•
Elements consisted of tiny "indestructible" particles called atoms.
•
Atoms of different elements have unique sizes and properties.
– The reason an element is pure is because all atoms of an element were
identical and that in particular they had the same mass.
– He also said that the reason elements differed from one another was that
atoms of each element were different from one another; in particular, they
had different masses.
– An atom of one element can't be change to an atom on another element.
•
Compounds are made of atoms of different elements combined
together.
– Compounds are pure substance because the atoms of different elements
are bonded to one another and are not easily separated from one another.
– Compounds have constant composition because they contain a fixed ratio
of atoms and each atom has its own characteristic weight, thus fixing the
weight ratio of one element to the other.
– In addition he said that chemical reactions involved the rearrangement of
combinations of those atoms.