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:
forms over the equator, where the rays of sunlight are the strongest and flow toward the poles, as well as in the stratosphere;
forms when oxygen gas (02) is exposed to ultraviolet radiation in the upper stratosphere
most of ozone is stored in the lower stratosphere
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 -
organic: carbon-containing chemicals
inorganic: non-carbon containing chemicals
physical: behavior and changes of matter
analytical: components and composition of substances
biochemistry: matter and processes of living organisms
environmental: matter and the environment
. . .
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:
time - second (s)
length - meter (m)
mass - kilogram (kg)
temperature - Kelvin (K)
amount of a substance - mole (mol)
electric current - ampere (A)
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:
Tera
Giga (G) - 109
Mega (M) - 106
Kilo (K) - 103
Hecto (H) - 102
Deka (D) - 101
base - 100
deci (d) - 10-1
centi (c) - 10-2
milli (m) - 10-3
micro (µ) - 10-6
nano (n) - 10-9
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:
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)
multiplication: multiply the coefficients and add the exponents
division: divide the coefficients and subtract the exponents
V. Significant Figures
rules for significant figures -
nonzero digits are always significant
zeros between nonzero numbers are always significant
all final zeroes to the right of the decimal point are significant
placeholder zeros are not significant - to remove them, write the number in scientific notation
counting numbers and defined constants have an infinite number of significant figures
significant figures: the digits reported in an answer; rules -
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
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 -
if the digit to the right of the last significant figure is less than 5, do not change the last significant figure
if the digit to the right of the last significant figure is greater than 5, round up the lsat significant figure
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
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 -
identify the independent and dependent variables
determine the range of the data that needs to be plotted for each axis
choose the ranges for the axes
number and label each axis
plot the data points
draw the “best fit” line
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 -
solid: has its own definite shape and volume with strong attraction force- “vibrates”
liquid: has a definite volume but takes the shape of the container with weak attraction force - “flows”
gas: has no definite shape or volume with very weak attraction force ; expands to fill a container - expands apart”
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
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 -
matter is composed of extremely small particles called atoms
atoms are indivisible and indestructible
atoms of a given element are identical in size, mass, and chemical properties
atoms of a specific element are different from those of another element
different atoms combine in simple whole number ratios to form compounds=
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 Na11- 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 - 32O162- -> 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; 0e1-; 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 -> oe1- + 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 + 0e1- -> 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) -
radio waves
cell phone waves
microwaves
radar waves
infrared waves
visible light waves
ultraviolet light
X-rays
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 -
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
l, angular quantum number: tells the shape of the orbital; for orbitals of the same principal energy level, l distinguishes different shapes
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)
ms 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 -
s: contains 1 orbital, 2 electrons; spherical shaped, therefore 1 possible orientation
p: contains 3 orbitals, 6 electrons; dumbbell shaped
d: contains 5 orbitals, 10 electrons; 2 dumbbell shaped
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 -
Aufbau principle: states that an electron occupies the lowest energy orbital in order of increasing energy
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
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 -
find the element on the periodic table
go up a period (row) and put the symbol of the noble gas in brackets
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 -
he published the 1st periodic table, that was arranged so that elements in the same group have similar properties
elements were arranged in order of increasing atomic mass
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