chem unit 1-3

Electromagnetic Spectrum

Electromagnetic Spectrum

The range of all possible frequencies of electromagnetic radiation.

Wavelength

The distance between successive crests of a wave, commonly denoted by the symbol λ.

Frequency

The number of crests of a wave that move past a given point in a given unit of time.

Energy of a Wave

The amount of energy carried by a wave, directly proportional to its frequency.

Speed of Light Equation (c = λν)

The equation relating the speed of light (c), wavelength (λ), and frequency (ν) of a wave.

Atomic Emission Line Spectrum

The spectrum of light emitted when an electron returns to a lower energy level.

Continuous Spectrum

A spectrum that contains all wavelengths within a given range.

Line Spectrum

A spectrum showing only certain discrete wavelengths.

Heisenberg Uncertainty Principle

States that it is impossible to know both the exact position and momentum of a particle simultaneously.

Orbital

A region in an atom where there is a high probability of finding electrons.

Aufbau Principle

States that electrons fill orbitals starting at the lowest available energy state before filling higher states.

Pauli Exclusion Principle

States that no two electrons in an atom can have the same set of four quantum numbers.

Hund's Rule

States that electrons will occupy empty orbitals of the same energy level before pairing up.

Orbital Diagrams

A visual representation of the arrangement of electrons within orbitals of an atom.

Electron Configuration

The distribution of electrons of an atom in its atomic orbitals.

Periodic Table

A tabular arrangement of the chemical elements, organized by atomic number and electron configuration.

effective nuclear charge (Zeff)

amount of charge felt by the most recently added electron

shielding effect

the reduction of the attractive force between a nucleus and its outer electrons due to the blocking effect of inner electrons

atomic radius

size of an atom

ioniation energy

the energy required to remove an electron from an atom

Reactivity

The ease and speed with which an element combines, or reacts, with other elements and compounds.

nuclear charge increases as you move ______ and ____ along a period

up and right

shielding effect increases as you move ____ and ____ along a period

down, left

Atomic radius increases as you move ____ and ____ along a period

down and left

ionization energy decreases as you move ____ and increases _____ along a period

down, left

reactivity trend in nonmetals

decreases down a group, increases across a period

Periodic Law:

States that when elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

Groups:

Vertical columns of element that have similar physical and chemical properties.

Period:

Horizontal row of elements in the periodic table.

Metals:

-Elements that are good conductors of heat and electrical current.

-They are solid at room temperature

-Most are malleable

-Many are ductile

Non-metals:

-Elements that tend to be a poor conductor of heat and electrical current.

-Brittle

-They have opposite properties of metals

Metalloids:

-Elements that tend to have similar properties of metals and non-metals.

-All solids

-semiconductors of heat and electricity

-Can be changed by using mixtures (ex: pure silicon is a poor conductor of electricity but when mixed with boron, it produces a good electrical conductor).

From which area on the periodic table are you most likely to find semiconductors/metalloids?

They are found along the staircase

Modern Periodic table:

-118 elements

-18 groups

-7 periods

-Arranged by atomic number

-A predictive model based on the arrangements of electrons in the elements.

Main Group Elements:

-Also called representative elements, are elements in Groups 1, 2, and 13-18.

-They display a wide range of physical and chemical properties.

-In their atoms, the s and p sublevels in the highest occupied energy levels are partially filled.

Alkali Metals:

-Group 1

-Very reactive

-Good conductors

-Malleable

-Ductile

-One valence electron

-Can explode when exposed to water

-They want to dump off an electron to become a positive ion, cation +1

Alkaline Earth Metals:

-Group 2

-Low electron affinities

-Two electrons in the outer shell

-Smaller atomic radii than Alkali Metals

-form cations +2

Halogens:

-Group 17

-Extremely reactive gases that form anions, ions with one negative charge

-Seven valence electrons

-Highly reactive with alkaline earth and alkali metals

Noble Gases:

-Group 18

-Fairly nonreactive

-Complete valence shell

-All gases at room temperature

Electron Configuration:

-The way electrons are arranged around the nucleus of an atom. An element's chemical and physical properties are due to the element's electron configuration.

-You can use the periodic table to help determine an element's electron configuration and its properties.

-Elements in the same group of the table have related outer shell configurations.

Core electrons:

Inside electrons. To find the number of core electrons you subtract the number of valence electrons from the total amount of electrons in the atom.

What does Coulomb's law determine?

Magnitude of force between two charges

How does Coulomb's law help us interpret interactions in an atom?

By showing how strongly the nucleus attracts different electrons

Shielding Effect:

-the core and valence electrons are toxic and valence keeps wanting to go in the nucleus but core pushes the valence away

-core and valence electrons make out with each other basically

-The amount of shielding an electron experiences depends on which orbital the electron is in. Inner core electrons always shield valence electrons, but valence electrons don't shield one another.

Effective Nuclear Charge, Zeff:

-increases left to right on the periodic table

-decreases down a group.

Silicon and germanium are in the same group. How does the effective nuclear charge of silicon compare to that of germanium? Why is this the case?

They have the same effective nuclear charge because they have the same number of valence electrons. Silicon does, although, have a greater force because it has fewer shielding electrons.

Atomic Radius:

-One-half the distance between the nuclei of the two atoms of the same elements.

-Increases right to left

-increases down a group

Atomic Radius in groups and periods:

-In groups the atomic radius increases down a group because valence electrons occupy higher energy shells

-In periods, the atomic radius decreases across a period because the number of electrons are changing.

Cations:

-Cations have a positive charge of 1+

-Metals tend to lose electrons to become cations.

-Cations are smaller than their parent ions

-The charge of a cation depends on the number of valence electrons that have been lost.

Anions:

-Anions have a negative charge of 1-

-Non-metals tend to gain electrons to become anions.

-Anions are bigger than their parent ions

Why are cations always smaller than their parent atoms?

Cations are always smaller than their parent ions because they're losing an electron. When they lose electrons they loose energy shells, creating a smaller atomic radius.

Why are anions always bigger than their parent atoms?

Anions are always bigger than their parent ions because they're gaining electrons. Since the number of protons stay the same, there is a lower Zeff per electron, which allows the extra electrons move farther from the nucleus.

Ionization Energy:

-The energy required to remove the outermost electron from the ground state

-Increases left to right and increases top to bottom

-The size of an atom directly affects its ionization energy.

-Small atoms tend to have high ionization energy and large atoms have lower ionization energies.