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

DefinitionsModule 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.

Dalton's Atomic Model

1800

1900

Rutherford's' Model

1920

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 ~1939The Structure of an Atom

• Each atom is made up of:

– Protons and Neutrons in the nucleus

– Electrons on the outside of the atom

Bohr ModelEnergy 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)

– Beta decay - Neutron -> proton + electron (beta particle) energy– Example U239(92P) -> NP239(93P)

• Alpha decay - 2N + 2 P (He nucleus leaves the atom) + energy– Example Po 214 (84P) -> Pb 210 (82P) + alpha particle (He nucleus) +E– See Fig13.56 page 328

• Gamma decay – 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)– 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 C12Half life is 5700 yearsAssumption 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 – H2O 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 Pure Substances

Heterogeneous Homogeneous 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.