Chemistry Midterm Study Guide

I.  The Ozone Layer

• chemistry: the study of everything around us; the study of the composition, structure, and properties of matter and the changes it undergoes

• matter: building blocks of everything in the universe and everything made from them; anything that occupies space and has mass

• ultraviolet light damages living organisms (animals and plants)

• troposphere: the lowest level of the atmosphere made up of the air we breathe

• stratosphere: the second lowest layer of the atmosphere that contains the protective ozone layer 

• ozone layer: layer of ozone in the stratosphere that absorbs ultraviolet light, protecting earth’s organisms

• ozone: 

1. forms over the equator, where the rays of sunlight are the strongest and flow toward the poles, as well as in the stratosphere;

2. forms when oxygen gas (02) is exposed to ultraviolet radiation in the upper stratosphere

3. most of ozone is stored in the lower stratosphere

4. can be measured by grounded instruments or balloons, satellites, and rockets

• ammonia (NH3): a tox gas in the form of coolants which escape from refrigerators and harms the ozone layer

• chlorofluorocarbons (CFCs): used as coolants in refrigerators as a replacement for ammonia; they are synthetic (made in a laboratory), stable (don’t react with other chemicals), and nontoxic; made up of carbon, fluorine, and chlorine

II.  Mass, Weight, and Models 

• mass: measurement that reflects the amount of matter; is constant and never changes, not dependant on gravity (SI unit - kilogram/kg)

• weight: measure of mass and the force of gravity on an object; can change from place to place; dependant on gravity; (c x mass or mass x gravity)

• if gravity increases, so does weight 

• much of matter and its behavior can be described at macroscopic; matter’s structure, composition, and behavior can be described a submicroscopic

• macroscopic: described easily with eye vision without a microscope

• submicroscopic/microscopic/atomic: more in depth observation, requiring a microscope

• chemistry explains events on the submicroscopic level that cause macroscopic observations, leading to the formation of models 

• model: verbal, visual, or mathematical explanation of experimental data

• scientific model: built on investigation and observation and explain macroscopic observations

• branches of chemistry - 

1. organic: carbon-containing chemicals

2. inorganic: non-carbon containing chemicals

3. physical: behavior and changes of matter

4. analytical: components and composition of substances

5. biochemistry: matter and processes of living organisms

6. environmental: matter and the environment

7. . . .

III.  Data Relationships 

• scientific method: the systematic, methodical backbone to every science experiment; involves repeated steps until a hypothesis is supported or discarded

• observations -> hypothesis -> experiments/conclusions/revised hypothesis -> theory/experiments/revised theory -> theory or scientific law

• quantitative data: data involving numbers followed by their units

• qualitative data: data involving physical characteristics (color, odor, shape, measurements, . . .)

• observations: quantitative and qualitative data combined

• hypothesis: an educated guess on your findings and observations always beginning with “If”

• experiment: the tested procedure

• conclusion: final findings from the experiment leading to form a revised hypothesis; states if you accept or deny your original hypothesis

• revised hypothesis: second hypothesis made if you deny your original hypothesis, formed off of your conclusion 

• scatter plots: show how a change in 1 variable influences another variable

• direct relationship: when 1 variable increases and so does the other

• no relationship: points appear to be a collection of dots; no clear involvement between the variables

• inverse relationship: when 1 variable increases, the other decreases

• theory: explanation that has been repeatedly supported by many experiments (ie Atomic Theory, Collision Theory of Reactions, Theory of Gravity)

• scientific law: relationship in nature that is supported by many experiments with no exceptions to these relationships being found; directly states what happens (ie Law of Gravity, Laws of Motion)

• independent variable: variable in an experiment that begins the data set and influences the dependent variable 

• dependent variable: variable in an experiment that changes based off of the independent variable

• constant: materials in an experiment that remain the exact same throughout the entire experiment

IV.  Research and Discovery

• basic/pure/fundamental research: research undertaken to gain knowledge for the sake of knowledge itself; leads us to a better understanding of how the natural world operates

• applied research: research undertaken to solve a specific problem; aims to develop useful applications from the knowledge gained from basic research

• chance discoveries: discovers formed when scientists obtain results far different than what was expected; happen spontaneously; used to improve and convenience our lives 

CHAPTER 2

I.  Base Units 

• quantity: something that has magnitude, size, or amount

• unit: a quantity adopted as a standard of measurement

• base unit: a defined unit in a system of measurement that is based on an object or event in the physical world, and is independent of other units 

• SI (Standard International) units:

1. time - second (s)

2. length - meter (m)

3. mass - kilogram (kg)

4. temperature - Kelvin (K)

5. amount of a substance - mole (mol)

6. electric current - ampere (A)

7. luminous intensity - candela (cd)

• Zero Kelvin: the point where there is virtually no particle motion or kinetic energy, also known as absolute zero 

• temperature: a measure of the average kinetic energy of the particles in a sample; heat transfers from hot -> cold

• 0° C = 32° F (freezing point)

• 100° C = 212° F (boiling point)

• 1°C = 9/5° F

• 37°C = 98.6° F

• K = C + 273.15

• F = 1.8 x C + 32

II.  Conversions and Metric Ladder

• meter: the starting point for the metric system

• metric ladder:

1. Tera 

2. Giga (G) - 10⁹

3. Mega (M) - 10⁶

4. Kilo (K) - 10³

5. Hecto (H) - 10²

6. Deka (D) - 10¹

7. base - 10⁰

8. deci (d) - 10^-1

9. centi (c) - 10^-2

10. milli (m) - 10^-3

11. micro (µ) - 10^-6

12. nano (n) - 10^-9

13. pico (p) - 10^-12

III.  Density

• measurement: a quantitative observation consisting of 2 parts

• derived unit: unit that is defined by a combination of base units (ie area and volume)

• 1 mL = 1 cm³

• 1 L = 1 dm³

• volume’s SI derived unit is cm³

• density: a physical property that measures the amount of mass per unit volume (derived unit: g/cm³); all substances have density - liquids, solids, and gasses

• the more dense substance will be on the bottom, while a less dense one will float to the top

• density formula: D = mass/volume

• density triangle:

            M

—------------------

      D    ||      V  

• to determine the density of an irregular shape, multiply its mass by the displacement of a liquid by that irregularly shaped object (the volume of the irregular object would be equal to the difference of the water level before the object entered it to after it entered)

IV.  Scientific Notation 

• scientific notation: 

1. addition and subtraction: write each quantity with the same exponent, then add or subtract the coefficients (ie 3.0 x 10⁶ + 4.0 x 10⁶ = 7 x 10⁶)

2. multiplication: multiply the coefficients and add the exponents

3. division: divide the coefficients and subtract the exponents

V.  Significant Figures

• rules for significant figures -

1. nonzero digits are always significant

2. zeros between nonzero numbers are always significant

3. all final zeroes to the right of the decimal point are significant

4. placeholder zeros are not significant - to remove them, write the number in scientific notation

5. counting numbers and defined constants have an infinite number of significant figures 

• significant figures: the digits reported in an answer; rules - 

1. addition and subtraction: the answer cannot have more digits to the right of the decimal point that any of the original numbers - round to that digit

2. multiplication and division: the number of significant figures in the answer is set by the original number that has the smallest number of significant figures 

• rounding numbers -

1. if the digit to the right of the last significant figure is less than 5, do not change the last significant figure

2. if the digit to the right of the last significant figure is greater than 5, round up the lsat significant figure

3. if the digits to the right of the last significant figure are a 5 followed by a nonzero digit, round up to the last significant figure

4. if the digits to the right of the last significant figure are as followed by a 0 or no other number at all, look at the last significant figure; if it is odd, round it up - if it is even, do not round up

VI.  Accuracy and Precision

• accuracy: how close a measured value is to an accepted value or the true value

• precision: how close a series of measurements are to one another

• error: the difference between the experimental value and the accepted value; |experimental value - accepted value|

• experimental value: the value measured in the lab

• accepted value: the correct value based on reliable references 

• percent error: expresses error as the percentage of the accepted value; indicates accuracy of a measurement; |error|/accepted value x 100

VII.  Graphing

• graphs: useful tools in science; the visual characteristics of a graph make trends in data easy to see

• steps in setting up a graph - 

1. identify the independent and dependent variables

2. determine the range of the data that needs to be plotted for each axis

3. choose the ranges for the axes

4. number and label each axis

5. plot the data points

6. draw the “best fit” line

7. give the graph a title

• directly proportional: as the independent variable (x axis) increases, the dependent variable (y axis) increases as well

• inversely proportional: as the independent variable increases, the dependent variable decreases

• interpolation: reading and estimating values falling between points on the graph

• extrapolation: estimating values outside the points by extending the line

• bar graphs: used to show how a quantity varies across categories

• circle graph/pie chart: has wedges that visually represent percentages of a fixed whole; shows the amount each part makes up of the whole

• slope: the ratio of the vertical change to the horizontal change; y₂-y₁/x₂-x₁

CHAPTER 3

I.  States of Matter

• matter: anything that has mass, volume, and takes up space; everything around us

• substance: a form of matter with a uniform and unchanging composition

• states of matter - 

1. solid: has its own definite shape and volume with strong attraction force- “vibrates”

2. liquid: has a definite volume but takes the shape of the container with weak attraction force - “flows”

3. gas: has no definite shape or volume with very weak attraction force ; expands to fill a container - expands apart”

4. plasma: extraordinary state of matter, consisting of high energy particles - have an indefinite shape and volume, like gasses; electrons are stripped from their nuclei; the most abundant and plentiful state of matter in the universe; is also ionized gas and a very good conductor of electricity, affected by magnetic fields (ie fluorescent light, stars, lightning, flames

5. Bose-Einstein: the “newest” state of matter that is harder than solids with a definite shape and volume; exists at extremely cold temperatures; particles are very unexcited, clumping together so firmly that they move as a single unit 

• melting: solid to liquid

• vaporization: liquid to gas

• ionization: gas to plasma

• recombination: plasma to gas

• condensation: gas to liquid

• freezing: liquid to solid

• sublimation: solid to gas; skips the liquid phase

• deposition: gas to solid; skips the liquid phase

II.  Properties of Matter

• intensive properties: independent of the amount of the substance that is present (ie density, boiling point, color, conductivity)

• extensive property: depends upon the amount of the substance present (ie mass, volume, energy)

• physical properties: can be observed without changing a substance into another substance or its chemical composition (ie boiling point, density, mass, volume, color, solubility)

• chemical properties: can only be observed when a substance is changed into another substance (ie flammability, corrosiveness/rusting, reactivity with acid, combustibility) 

III.  Changes of Matter

• physical change: does not alter the composition or identity of a substance (ie ice melting, sugar dissolving in water)

• chemical change: alters the composition or identity of the substances - cannot be observed without changing the composition of a substance; one or more substances turning into a new substance involved (ie hydrogen burns in air to form water, decomposing, rusting, exploding, burning, oxidizing)

• subscripts: tell how many of a particular type of an atom are inside a molecule

• coefficients: tell how many of each particle is involved in the reaction

IV.  Types of Matter

• mixture: a type of matter that can be separated physically

• pure substance: a type of matter that cannot be separated physically

• heterogenous mixture: a type of mixture that’s differences are unevenly mixed; separate parts can be seen (ie dirt, milk, blood)

• homogeneous mixture/solution: a type of mixture that is uniform and evenly mixed; looks the same throughout (ie lemonade, gasoline, steel)

• element: the simplest form of matter that cannot be separated physically or chemically; is only 1 type of ato and retains its properties

• compound: different elements chemically bonded together; can be chemically, not physically, separated; a completely new substance with new properties (ie carbon dioxide CO₂, water H₂O, salt NaCl)

V.  Methods to Separate Mixtures 

• mechanical separation: often by hand - takes advantage of physical properties such as color and shape (ie recycling plastic, paper, and metal, sifting or sieving, hand separation)

• magnet

• filtration: takes advantage of the physical property of the state of matter; a screen lets the liquid particles through, but traps the solid particles (ie filtering coffee, spaghetti)

• evaporation: vaporizing a liquid and leaving the dissolved liquids behind; used to separate solution (ie obtaining sea salt from sea water evaporation ponds)

• decanting: to pour off a liquid, leaving another liquid or solid behind; takes advantage of differences in density

• extraction: filtration + evaporation - used to separate an insoluble solid (something that does not dissolve in a liquid) from a soluble solid (something that does dissolve in a liquid; done by adding a solvent (liquid that does the dissolving) to the mixture, then pouring the liquid through a filter

• distillation: the separation of a mixture of liquids based on the physical property of boiling point; takes advantage of different boiling points (ie the distillation of alcohol or oil)

• density separation: more dense components sink to the bottom and less dense components float 

• fractional crystallization: dissolved substances crystallize out of a solution once their solubility limit is reached as the solution cools (ie growing rock candy or the crystallization of a magma chamber)

• paper chromatography: uses the property of a molecular attraction (molecular polarity) to separate a mixture; different molecules have varying molecular attractions for the paper (the stationary phase) vs the solvent (the mobile phase) (ie the separation of plant pigments and dyes)

• centrifuge: circular motion helps denser components sink to the bottom faster (ie the separation of blood or DNA from blood)

• sublimation: the process of a solid changing directly to a gas, which can be used to separate mixtures of solids when 1 sublimates and the other does not 

VI.  Periodic Table

• periodic table: organizes the elements into a grid of of periods and groups

• periods: horizontal rows on the periodic table

• groups: vertical columns on the periodic table

• metals: found on the left side of the staircase on the periodic table (except hydrogen); have few electrons in their outer energy level, thus losing electrons easily; ductile, good conductors, malleable, lustrous, most are solid at room temperature (ie sodium Na, gold Au)

• nonmetals: found on the left side of the staircase; most have almost full outer energy levels, thus tending to gain electrons (some have completely full outer levels of 8 electrons); not ductile, not malleable, not lustrous, poor conductors of electricity, most are solid, some are gas 

• metalloids: border the staircase; most atoms have about half a complete set of electrons in their outer levels (4 electrons); have properties of both metals and nonmetals 

VII.  Laws

• law of conservation of mass: states that mass is neither created nor destroyed in chemical reaction - it is conserved from the reactants to the product; mass of the reactants equals the mass of the products

• law of definite proportions: states that a given compound always contains elements in a certain proportion by mass; a compound is unique because of the specific arrangement and weights of the elements which make that compound 

• mass percentage of an element (%) = (mass of element/mass of compound) x 100

• law of multiple proportions: states that when 2 elements form 2 different compounds, the masses of 1 element that combine with 1 gram of the other element can be expressed as a ratio of small whole numbers; same elements, different compounds with different ratios of elements; different compounds can be made with different ratios of the same elements

CHAPTER 4 

I.  Dalton’s Atomic Theory

• Dalton’s Atomic Theory: proposed in 1803; key points - 

1. matter is composed of extremely small particles called atoms

2. atoms are indivisible and indestructible

3. atoms of a given element are identical in size, mass, and chemical properties

4. atoms of a specific element are different from those of another element

5. different atoms combine in simple whole number ratios to form compounds=

6. in a chemical reaction, atoms are separated, combined, or rearranged 

II.  Atoms, Ions, and Isotopes

• proton (p⁺): a positively charged subatomic particle located in the nucleus of an atom; weighs 1 amu (atomic mass unit) and has a charge of 1+; the number of protons equals the number of electrons in a neutral atom

• electron (e^-): a negatively charged subatomic particle located in the electron cloud, the outer energy level of an atom; has no weight, weighing 0 amu, and has a charge of 1-; the number of electrons equals the number of protons in a neutral atom

• neutron (n⁰): a neutrally charged subatomic particle located in the nucleus of an atom; weighs 1 amu and has a neutral charge of 0

• atomic number: the number of protons in an atom; the smaller, whole number in an element’s chemical symbol 

• atomic mass/mass number: the number of protons plus the number of neutrons in an atom, located in the nucleus; the larger, often decimal number in an element’s chemical symbol

• ion: a positively or negatively charged atom, caused by a change in its number of electrons (ie Na^11- is now a negatively charged ion, because it has gained 11 electrons); when an atom gains electrons, it becomes negatively charged, because there are more electrons than protons - when an atom loses electrons, it becomes positively charged, because there are more protons than electrons

• cation: a positively charged ion (has lost electrons)

• anion: a negatively charged ion (has gained electrons)

• isotope: an atom of a particular element that has a different atomic mass than other atoms of that element, caused by a change in its number of neutrons in the nucleus; despite having different masses, they do not differ significantly in chemical behavior (to determine if an atom is an isotope, see if its atomic mass is different from the atomic mass of that element listed on the periodic table)

• example -  ³²O₁₆^2- -> Oxygen, 16 protons, 18 electrons, 16 neutrons, ion, anion, atomic number of 16, atomic mass of 32

• average atomic mass: the weighted average of all the naturally occurring isotopes of a particular element; depends on both the mass and the relative abundance of each of the element’s isotopes that are used in the equation; formula: (mass x %) + (mass x %) / 100

III.  Nuclear Chemistry

• nuclear reactions: involve changes in atomic nuclei

• spontaneously-changing nuclei emit radiation, and are said to be radioactive

• radioisotopes: isotopes of atoms with unstable nuclei, spontaneously releasing protons and neutrons from the nucleus -  are radioactive and unstable

• stable isotopes: isotopes of atoms that do not release protons or neutrons from the nucleus are not radioactive (normal isotopes)

• nucleon: a proton or a neutron

• nuclide: a nucleus with a specified number of protons and neutrons

• radionuclides: radioactive nuclides, spontaneously emitting radiation - are radioactive and unstable

• alpha particle: 𝛂; ⁴He₂²⁺; heavy mass

• alpha particle emission/alpha decay: decreases the atomic mass by 4 and the atomic number by 2; low penetration and danger, protected by skin (ie ²³⁴U₉₂ -> ⁴He₂ + ²³⁰Th₉₀)

• beta particle: 𝛃; ⁰e_1-; light mass

• beta particle emission/beta decay: converts a neutron into a proton, increasing the atomic number by 1; medium penetration and danger, protected by paper or clothing (ie ²³⁴Pa₉₁ -> ^oe_1- + ²³⁴U₉₂) 

• gamma ray: 𝛄; ⁰𝛄₀; no mass

• gamma ray emission/gamma radiation: does not change the nucleus or the chemical symbol; high penetration and danger, protected by lead; consists of high energy protons and is emitted when nucleons rearrange into a more stable configuration in the nucleus; often accompanies other nuclear decays, and does not turn an element into another

• positron decay: converts a proton into a neutron, decreasing the atomic number by 1 (ie ²³Mg₁₂ -> ⁰e₁₊ + ²³Na₁₁) 

• positron: the antiparticle of the electron

• neutrino: a massless, chargeless particle (ie ⁰𝛄₀)

• electron capture: an inner-orbital electron is captured by the nucleus; converts a proton into a neutron (ie ¹¹C₆ + ⁰e_1- -> ¹¹Be₅)

• not all combinations of protons and neutrons (in the nucleus) are stable; to determine stability, N/Z is examined, where N=neutrons and Z=protons

CHAPTER 5 

I.  Wave Nature of Light

• electrons spread through space as an energy wave

• light travels through space in the form of radiant energy; it travels through waves or as fast-moving particles

• visible light: a type of electromagnetic radiation

• electromagnetic radiation: a form of energy that exhibits wave-like behavior as it travels through space

• amplitude: how much energy a wave carries; the more energy, the higher the amplitude; the vertical distance from the origin to the crest 

• wavelength: the distance between a point on a wave and the nearest point just like it (measured in m, nm; 1 m = 10⁹ nm)

• frequency: the number of wavelengths that pass a fixed point each second; as frequency increases, energy increases and wavelength decreases (measured in Hz or s^-1)

• crests: the highest points of a wave

• troughs: the lowest points of a wave

• hertz: 1 wavelength passing by a fixed point per second; measures frequency; equal to 1 s^-1

• electromagnetic spectrum (from lowest to highest frequency) - 

1. radio waves

2. cell phone waves

3. microwaves

4. radar waves

5. infrared waves 

6. visible light waves 

7. ultraviolet light

8. X-rays

9. gamma rays

• continuous spectra: contains all colors (R O Y G B I V)

• line spectra: sharp loans observed in a spectrum of light emitted or absorbed by an element 

• flame tests: many elements give off characteristic light which can be used to help identify them (atoms are excited by heat or electricity)

• c = λν ; speed of light = wavelength x frequency

• speed of light: 3.00 x 10⁸ m/s

• the speed for all electromagnetic waves is the speed of light in a vacuum, 3.00 x 10⁸ m/s

• E = hν ; energy = Planck’s constant x frequency

• Planck’s constant: 6.626 x 10^-34 J x s

• if you need to find E, energy, but the frequency is not given, use the formula - E = h x (c/λ) 

II.  Particle Nature of Light

• quantum: the minimum amount of energy that can be gained or lost by an atom; matter gains or loses energy only in small, specific amounts, quanta

III.  Quantum Mechanical Model of the Atom

• ground state: the state at which an atom has its lowest energy level

• excited state: state of an atom caused by the excitation of the electron by absorbing energy

• when atoms absorb energy -> electrons move onto higher atomic levels; these electrons then lose energy by emitting light when they return to lower energy levels

• principal energy levels: the regions of space in which electrons can move about the nucleus

• Heisenberg uncertainty principle: states that is is fundamentally impossible to know precisely both the velocity and position of a particle at the same time; you can determine where the electron is (position/orientation), but not where it is going (speed/momentum), and vise versa

• de Broglie equation: predicts that all moving particles have wave characteristics; λ = h/mv (wavelength = Planck’s constant/ mass of the particle x velocity)

• atomic orbital: a region of space in which there is a high probability of finding an electron; each orbital contains a maximum of 2 electrons

• quantum numbers: specify the address of each electron in an atom and describe the properties and the electrons in orbitals - 

1. n, principal quantum number: electron’s energy depends principally on this (tells which level the electron is in, ie n=1, n=2, . . . ); are positive numbers beginning with 1 (closest to the nucleus and moving out); as n increases, so does the energy of the electron

2. l, angular quantum number: tells the shape of the orbital; for orbitals of the same principal energy level, l distinguishes different shapes 

3. m₁, magnetic quantum number: indicates the orientation of an orbital around the nucleus; s orbital is sphere shaped, p orbital is shaped dumbbell shaped, d orbital is like 2 dumbbells, and f orbital is flower shaped)

4. m_s spin quantum number: has 2 possible values (clockwise ½ or counterclockwise -½), which indicate the 2 fundamental spin states of an electron in an orbital; identifies the 2 possible spin orientations of an electron in an orbital (clockwise ½ or counterclockwise -½)

• 2n²: formula for the number of electrons that can fit in a shell/principal energy level (ie n=3, 2 x 3² = 18; 18 electrons can fit and be contained in n=3)

• quantum theory: describes mathematically the wave properties of electrons and other very small particles

• sublevels/energy sublevels: are contained within the principal energy levels (rooms inside the main floors, the principal energy levels); include s, p, d, and f -

1. s: contains 1 orbital, 2 electrons; spherical shaped, therefore 1 possible orientation

2. p: contains 3 orbitals, 6 electrons; dumbbell shaped

3. d: contains 5 orbitals, 10 electrons; 2 dumbbell shaped

4. f: contains 7 orbitals, 14 electrons; flower shaped

IV.  Electron Configuration and Orbital Notation

• electron configuration: tells us in which orbital the electrons for an element are located, following 3 rules - 

1. Aufbau principle: states that an electron occupies the lowest energy orbital in order of increasing energy

2. Pauli exclusion principle: states that a maximum of 2 electrons can occupy a single orbital, but only if the electrons have opposite spins; no 2 electrons can fill 1 orbital with the same spin; also, no 2 electrons in the same atom can have the same set of 4 quantum numbers

3. Hund’s rule: states that orbitals of equal energy are each occupied by 1 electron before any orbital is occupied by a second electron, and all electrons in sumpy occupied orbitals must have the same spin

• order/chart for electron configuration - 

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

7s 7p 7d 7f

• example (Nitrogen, 7 e^-) - 

1s² 2s² 2p³

CHAPTER 6 

I.  Noble Gas Configuration

• noble gas configuration: shorter (abbreviated) way of finding ground electron configuration - 

1. find the element on the periodic table

2. go up a period (row) and put the symbol of the noble gas in brackets

3. finish the outer levels

II.  Element Classification

• metals: lustrous (shiny), solid at room temperature (except mercury), good conductors of heat and electricity, ductile (can be drawn into wire), malleable (can be hammered into sheets), high melting points, high density 

• nonmetals: to the right of the staircase on the periodic table; dull, many are gases at room temperature (except Bromine, liquid), poor conductors of heat and electricity, brittle, not ductile/malleable, low melting point, low density

• metalloids/semimetals: border the staircase on the periodic table; exhibit properties in between metals and nonmetals - dull and shiny, brittle, semi-conductors of heat and electricity

• alkali metals: group 1 (except Hydrogen); 1+ charge, very reactive

• alkaline earth metals: group 2; 2+ charge, less reactive than alkali metals

• halogens: group 17; 1- charge, very reactive

• noble gases: group 18; no charge, uncreactive (are stable because their outer energy level is full with 8 electrons)

• lanthanides: elements 58-71; contains f orbitals, part of f block

• actinides: elements 9–103; contains f orbitals, part of f block

• transition elements/transition metals: groups 3-12; variable charges

• main block/representative elements: groups 1, 2, 13-18

III.  Atomic Radius, Ionization Energy, Ionic Radius, Electronegativity

• atomic radius: measures the distance from the nucleus of 1 atom in a diatomic pair to the nucleus of the other atom, and divided by 2; commonly used unit is picometer, pm = 10^-12 m; decreases right on a period, increases down on a group

• diatomic pair: 2 atoms bonded together

• diatomic pair elements - I₂, Cl₂, H₂, N₂, Br₂, O₂, F₂ ; “I clearly have no brains or feet.”

• ionization energy: the amount of energy needed to remove the highest energy electron from an atom; increases right on a period, decreases down on a group 

• cation: positive ion, caused by the removal of electrons; metals readily lose electrons; always smaller than the neutral atom, caused by the loss of outer energy level electrons

• anion: negative ion, caused by the addition of electrons; nonmetals readily gain electrons; always bigger than the neutral atom, caused by gaining of electrons in their outer energy levels

• ionic radius: measures the distance from the nucleus of an ion to its outermost electron; the more electrons, the greater the ionic radius; if two ions have the same number of electrons, the ion with the greater number of protons (atomic number) will be smaller, because the more protons the greater the attraction force to the electrons, decreasing the ionic radius and the size of the ion

• electronegativity: the ability of an atom in a bond to pull on the electron or the tendency for a bonded atom to attract an electron to itself; when electrons are shared by 2 atoms, a covalent bond is formed -  when the atoms are the same, they pull on the electrons equally (ie H-H); when the atoms are different, the atoms pull on the electrons unevenly (ie HCI); increases right on a period, decreases down on a group

• F (fluorine) - most electronegative element 

• Cs (cesium) - least electronegative element

IV.  Development of the Periodic Table

• John Newlands: proposed the Law of Octaves and noticed that when the elements were arranged by increasing atomic mass, their properties repeated every 8th element; he arrange elements by atomic mass and noticed similar properties (56 known properties at this time)

• Dmitri Mendeleev: arranged the 63 known elements by increasing atomic mass and similar properties; he and Lothar Meyer each made a connection between atomic mass and properties of elements; they developed identically systems for identifying elements; Mendeelev received credit for his periodic table because - 

1. he published the 1st periodic table, that was arranged so that elements in the same group have similar properties

2. elements were arranged in order of increasing atomic mass

3. he left gaps on the table where he predicted future elements would belong

• Henry Mosely (1913): developed the atomic number (the number of protons in an atoms), becoming the basis for the modern day periodic table 

• periodic law: forms the basis for the organization of the Periodic Table; when elements are arranged in order of increasing atomic number, their physical and chemical properties show a predictable periodic pattern; the properties of yet undiscovered elements can be predicted based on their apparent location in the periodic table