5. Intro to Chemistry

• Chemistry

• Basics of Atoms and Molecules

• Reading Assignments• Text: AP Module 1 pages 1-20

• Homework Assignment• Module 1 Study Guide Questions p 24 # 1 -14

Introduction (p 1)

• Atoms and Molecules (p 1)

• Atom - the smallest chemical unit of a matter - very small

• - Contrast with a continuous view of matter• - Quick review of history of idea of atoms.• - As of 2007 117 different kind of atoms elements have

been observed as of 2007, • -- 94 occur naturally on Earth. • - Experiment 1.1 Atoms and Molecules (p 1)

• Molecules - Two or more atoms linked together to form a substance with unique properties. See figure 1.1 page 4. Examples

• - H2, O2, O3, H2O, H2O2

• Elements - Element, are pure substance that from form one type of atom that is defined by its atomic number; that is, by the number of protons in its nucleus. The term is also used to refer to a pure chemical substance composed of atoms with the same number of protons.

• Compounds - molecule made up of more than one type of atom that are chemically bonded together. The have unique properties that can not be broken down physically.

• Mixtures - substances mixed together physically but not chemically bonded.

Measurement and Units (p 7)

• Dimension is a measurable characteristic. Dimensions are used to describe the state or conditions of the physical world around us. Sometimes dimensions are referred to as a physical quality. Some examples of dimensions are length, mass, time, electric charge, area, speed, force, and weight.

• Fundamental dimensions describe the basic characteristics of the universe. Derived dimensions describe more complex characteristics of the universe that are made up of various combinations of fundamental dimensions.

• A unit is an agreed upon standard for measuring a dimension. It allows us to give a numerical value to a dimension. Sometimes dimensions are referred to as a physical quantity. Some examples of units are kilograms, feet, meters, seconds, minutes, square feet, square meters, miles per hours, meters per second and pounds.

• For example, mass (m) is a dimension and kilogram (kg) is a unit. Mass is the "measurable characteristic" being described by the "standard" or "unit" kilogram. To describe a quality of mass properly you need to have the number value and units such as 6 grams.

• PLEASE, ALWAYS REMEMBER that an answer is WRONG if the units are incorrect.

•

The Metric or Standard International System of Units

•

• There are many systems of units used to describe dimensions. In science the standard of units is called the International System of Units or SI system. This system is often called the metric system because many of the SI units were derived from the metric system. Some examples of SI units are kilograms, meters and seconds. The SI units are very logical in that they use prefixes based upon powers of 10 to describe large and small quantities. One tenth (1/10) of a meter is a decimeter and 1/100 of a meter is a centimeter. For example, a decameter is 10 meters; a hectometer is 100 meters.

• The other system of measurement we commonly use in the United States is called the English system. It uses units such as pounds, feet, and seconds. Since we use both systems in this country, we will use both in this course.

• See table 1.1 Page 9.– Mass -– Distance– Time– Volume– Weight Newton Slug Manipulation Units (p 10)

•

• See table 1.2

Table of Metric or SI Prefixes.

Converting Between United (p 11)

• Using the concept of multiplying by a form of "1" (P 12)

• Converting Between System (p 14)

• Again using the concept of multiplying by a form of "1" (P 14)– Experiment 1.2 Cubits and Fingers (p 15)

• Concentration (p 17)

• The quantity of a substance in a certain volume (somewhat like density)

• Experiment 1.3 Concentration (p 17)

The Forces in Creation Part 3 - Basics of Atoms

• Reading Assignments– Text: Module 13 pages 313 - 340

• Homework Assignment– Module 13 Study Guide Questions p 339-

340 # 1 -20

Introduction (p 313)

• You have learned about two fundamental forces (Gravitational and Electromagnetic)

• Now we will learn about the other two Strong and Weak Nuclear Forces

The Structure of an Atom (p313)

• Model: (See p 314) - A schematic description of a system that accounts for it known properties,

• More general: Some working representation of real systems

• Models are simplifications or idealizations of the real thing

• Some types of models:– - Parables– - Rules– - Diagrams or physical model like the globe– - Analogs– - Mathematical– - Computer

Atomic Models:

• All atomic models are mathematical models that can be represented by a diagram. Here are some examples of atomic models:

• Pre atomic: Matter is continuous• Greek: (Democritus 440 BC) - Matter is made up of discreet

fundamental particles that can't be divided. The also can only combine in certain ratios - Law of definite composition. H20, 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.

Dalton Model - Picture of Dalton's Theory of the atom

Thomson's Model (1890 - 1920) • Discovery of the electron experimenting with cathode (negative electrode) rays

Rutherford Model (1880's - 1930) • Discovery of the proton

Bohr Model• Energy Levels - Orbits• Electron's can only exist at certain energy levels or quanta• Photon (packet of light energy)• Beginning of quantum theory which explains very small - remember theory of

gravity can not explain the very small

Bohr Model Refined - The Proton - Neutron Model

• Chadwick Discovery: The Neutron: – Chadwick discovers the neutron as part of

the nuclease in 1932. Led to current module which is still in use.

Sub Atomic Particles

• Nucleus: Proton Neutron - weigh about the same -both about 1 amu Nutron sloghly heavier

• Electron: about 2000 times lighter than proton or Neutron

• Quarks make up protons, and Neutrons• Breakup of neutron into proton, electron and

antineutrino• Go over Bohr Model on p 314• Relative size of atom and parts - p 315 - electron size

baseball stadium - nuclus a marble

More on the Atom

• Atom 99.99% empty (p 316) • How do we "touch" • Interaction of electrons (p316) • Atomic number - number of protons -

dictates type of atom (p317)• Mass number (AMU) = Sum of Protons and

Nuetrons (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

• Current Quantum Mechanical Models (p 315 and page 323)

• Based upon the uncertainty principal and probabilities

• Bohr Orbits versus orbital, electron levels, or shells

• Orbit number and Electron capacity • 1 = 2, 2 = 8, 3 = 18, 4 = 32, 5 = 50 (see

table p 324)

Atoms, Elements, Compounds• An atom of one element can't be change to an atom on another

element.• Good place to contrast this as a hypothesis with the alchemist

hypothesis of Dalton's time. • 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.

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• 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 or transition metals B

• Periods or series are in rows– Metals– Metalloids– Nonmetals– Lanthanide Series - Actinide Series and result if they weren't modified

in shape (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)• Constant for metals

• Unusual things three letter, parenthesis recently discovered - rows that are below,

• Draw C12 Bohr Model - see page 323

• Draw Al 27 atom (p 324)

The Strong Force (p 325)

• Very strong short range force that keep the nucleus together P to P and N. (Note that as atoms get heavier the number of neutrons exceed the number of protons - provided more strong force to keep the Proton from repelling.

• Exchange of the pion particle causes the strong force

• Short lived particles - cause very strong forces - p 326

• A closer look at quarks - protons and neutrons made up of three quarks where the exchange of particle are called gluons

Radioactivity - the Weak Force (p 327)

• Radioactivity or radioactive decay Breakdown of an atom into teo or more atoms plus energy and particles are given off

• Caused by a release of energy when the weak nuclear force is released - similar to the EM force (now thought to be a different form of the same thing - like ice, liq after and vapor)

• Radioactive isotope: 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 nucluis leaves the atom) + energy • Example Po 214 (84P) -> Pb 210 (82P) + alpha particle (He nucleus)

+ E • See Fig13.56 page 328 • Gamma decay -Nuclus gives of high energy called gamma ray• 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 stem •

•

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

• 10 gram of U239(92P) -> 5 gram of U239(92P) = 5 grams of NP239(93P (see figu 13.6 page 333)

•

Radioactive Dating (p334)•

• Using the amount of radiaactive material in a substance to detrime 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.

Wonder of Water

• Reading Assignments

• Text: Module 4 pages 81 - 104

•

• Homework Assignment

• Module 4 Study Guide Questions p 104 # 1 - 14

Introduction (p 81)• Wonder of water• Look at its composition H2O - what would you think it would be

based upon its molecular weight?• Why:• Is it liquid at normal temperatures - needed for life• What can it change phase - weather - keeping balance temp• Why does it have such a high heat capacity for a simple

molecule• why does it expand when it freezes• Why can it hold more O2 when it gets colder 2• 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)•

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.•

Chemical Formulas (p 85)

•

• Using symbols to represent chemicals (H2O)

• Where names come from - some Latin ferrum so Fe for Iron.

•

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.

Division of Matter• Matter• Mixtures Pure Substances• Heterogeneous Homogeneous

Elements• Compounds• Elements and Their Symbols –

see periodic table

Chemical Bonds:• Sharing or “borrowing” outer shell – valence – electrons.• Follow rule of the octave• S - , P 8, D 8 and so on• Ionic bonds – borrowing electrons – not really consider a bond, but

an ionic attraction’note – electron with proton is intra-molecular interactions

• Intermolecular interaction - • Example Na+ Cl- • Covalent Bonds - sharing of electrons – true bond – very strong

bonds • Intermolecular bond• Single Bond• Double Bond• Triple Bond

Vader Walls Bonds• 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

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.

• Chemical compound: a substance consisting of two or more elements chemically-bonded together in a fixed proportion by mass. The basic unit (smallest unit that has these properties) of a compound is the molecule.

Chemical Symbols and Formulas of Compounds

• -- Use of subscript - goes with prior symbol

• -- Use of coefficient - in front of atom or compound

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.

• The following can indicate that a chemical change took place, although this evidence is not conclusive:

– * Change of color (e.g., rusting of iron causes a change in color from silver to reddish-brown).

– * Change in temperature or energy, such as the production (exothermic) or loss (endothermic) of heat.

– * Change of form (burning paper) (this change is difficult to reverse).– * An unexpected change in color– * Light, heat, or sound is given off.– * gasses formed, often appearing as bubbles.– * Formation of precipitate (insoluble particles).– * The decomposition of organic matter (rotting food)

• For example, placing a pot of water on a hot stove element causes a change in temperature and gas to be released (water vapor) but a chemical change did not take place. It was simply a physical change / change of state. An example could be a log that is burning.

• A chemical reaction produces new substances by changing the way in which atoms look. In a chemical reaction old bonds are broken and new bonds are formed between different atoms. This breaking and forming of bonds takes place when particles of the original materials collide with one another. An example of a chemical change is fireworks.