Unit 1 Quiz - H. Chemistry

John Dalton's Atomic Theory

1. All matter is made of tiny particles called atoms.

2. All atoms of a given element are identical.

3. Atoms of a given element are different from those of any other element.

4. Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative number and type of atoms.

5. Atoms are neither created nor destroyed in chemical reactions.

JJ Thomson: Cathode Ray Tube Experiment

• He put a stream of electrons into a electrified glass tube filled with gas where the left side was negatively charged and the right side was positively charged

• Electrons in the tube were attracted to the positive side

• This led to him discovering that electrons had a negative charge

  • led to plum pudding model

Thomson’s Plum Pudding Model

Thomson thought the atom was only electrons on a sphere of positive charge

Rutherford's Gold Foil Experiment

• took a source of alpha particles and shot a beam of them at a piece of gold foil surrounded by a sheet of zinc sulfide

• He found that most of the alpha particles went straight through but some bounced off of something and since electrons were too small to deflect the alpha particles he came to the conclusion that there was something else inside of the atom with a lot of mass

  • This led him to create the Nucleus Model of the atom where that something else was the nucleus

Rutherford: Nucleus Model

• that something else in the atom was the nucleus

  • much later scientists discovered that the nucleus is made up of protons and neutrons

• volume of the nucleus compared to the rest of the atom (the electron cloud) is actually quite small

• electrons are more likely to be found closer to the nucleus than further away

Electron (e^-)

• Charge of -1

• Found in the electron cloud

Proton (p⁺)

• Charge of +1

• Found in nucleus

• Determines the element

Neutron

• No Charge

• Found in nucleus

Nuclear Symbol Notation

• X = Element Symbol

  • The one to two letter abbreviation for the chemical element.

  • Has a superscript with its charge if it is charged

• A = Mass Number

  • The total number of protons + neutrons in the nucleus (written at the top left).

• Z = Atomic Number

  • The number of protons in the nucleus (written at the bottom left), which defines the element.

Periodic Table Notation

• X = Element Symbol

  • The one to two letter abbreviation for the chemical element.

• Z = Atomic Number

  • The number of protons in the nucleus, which defines the element.

• Atomic Mass

  • The weighted average mass of the element's atoms, usually measured in atomic mass units (amu).

Hyphen Notation

• X = Element Symbol

  • The one to two letter abbreviation for the chemical element.

  • Has a superscript with its charge if it is charged

• A = Mass Number

  • The total number of protons + neutrons in the nucleus (written at the top left).

Isotopes

Atoms with the same number of protons but different numbers of neutrons.

Average Atomic Mass (AAM)

• The weighted average mass of all the different isotopes of a particular element

  • Weighted by how abundant each isotope is

(review how to calculate with the lecture video)

Ions

• made by when a neutrally charged atom gains or loses electron(s)

• an atom with a charge

Anion

• Negatively charged ion

• Neutrally charged atom has gained electron(s)

• Nonmetals tend to gain one or more electrons to become anions.

(see notes of more info)

Cation

• Positively charged ion

• Neutrally charged atom has lost electron(s)

• Metals tend to lose one or more electrons to become cations.

Half-Life (t_1/2)

• The time required for a quantity to fall to half its value as measured at the beginning of the time period

• Used to describe a quantity undergoing exponential decay

• Is constant over the lifetime of the decaying quantity

• Radioactive atoms decay over time

  • man-made atoms typically have very short half-lives

• Radioactive nuclei with a shorter half-life is called “hotter” than one with a longer half-life

(review how to calculate with the lecture video)

Metals

Physical properties:

• Good conductors

• Malleable

• Ductile

• Lustrous

All elements to the left and below the “stair-step” line are metals except for hydrogen (H)

Non-metals

• Appear in the upper-right section of the periodic table

• Generally lack the properties that characterize metals

• More variation in properties than metals

  • Can be solids, liquids, or gases

Elements that border the staircase on the periodic table

Metalloids/Semisolids

Elements below the main section of the periodic table

Lanthanides and Actinides

Periodic table

an organized chart that shows all 118 known elements in order of increasing atomic number

Groups

• Elements with similar chemical behavior because each group contains elements that share the same number of valence electrons

• vertical columns

• 18 groups

Group 1

Alkali Metals

Group 2

Alkaline Earth Metals

Groups 3-12

Transition Metals

Group 17

Halogens

Group 18

Noble Gases

Period

• Each period represents a new principal energy level (electron shell) being filled with electrons.

• horizontal rows

• 7 periods

Electromagnetic radiation

It is a form of energy that acts like a particle and a wave, and its energy is defined by its wavelength and frequency.

wavelength (𝜆)

• the distance over which a wave's shape repeats, measured from one peak (crest) to the next peak, or from one low point (trough) to the next low point

• measured in units for length

Frequency (𝜈)

• number of waves that pass through a given point in a second

• measured in s^-1

Relationship between wavelength, frequency, and energy

• shorter wavelength, lower frequency, lower energy

• longer wavelength, higher frequency, higher energy

Visible Light

• The part of the electromagnetic spectrum that ranges from 400 nm to 700 nm.

  • Violet = 400 nm

  • Red = 700

photon

• packet of energy

• moves in streams

Emission of Energy by Atoms

Steps:

1. Atom absorbs energy

2. Electrons are promoted to excited state (electrons jump energy levels)

3. Atom is unstable

4. Atom releases energy in the form of a photon (energy absorbed = energy released)

5. Electrons return to ground state (electrons return to original energy levels)

Excited State

Higher energy state

Ground State

Lower energy state

Emission Spectra

the colors that are emitted when atoms release energy

different electron transitions (movement of an electron between energy levels) within an atom produce different colors (wavelengths) of light.

(Principal) Energy Levels/Shells

• Energy of shells increase the further they are from the nucleus

• Discrete energy levels = certain energy states that electrons can be at

• First shell: n=1, Second shell: n=2, etc.

• The jumps between energy levels are unique to the element because each element has a different number of protons in its nucleus, which affects how attracted the electrons are to the nucleus

• Ex: Subshells with an energy level of n=4

  • 4s, 4p, 4d, 4f

• Each energy level (or shell) can hold as many electrons as defined by 2n²

Bohr Model

• Very incorrect

• Electrons do not actually move in orbits the way planets move around the sum

• But easy to draw

The Wave Mechanical Model

• A "probability map" of finding the electron within a certain area; regions of space where the electron is "likely" to be found

  • Electrons are more likely to be found closer to the nucleus

• Orbitals

  • Each orbital can hold 2 electrons

• Schrödinger

Subshells/Sublevels

• Within a given energy level, there are sublevels (called subshells) (ex: s, p, d, f) which contain multiple orbitals of different shapes.

• In the energy level n=1 there is 1 subshell (1s), in the energy level n=2 there are 2 subshells (2s and 2p), in the energy level n=3 there are 3 subshells (3s, 3p, and 3d), etc.

s subshell

• Spherical subshell

• Each s subshell consists of 1 orbital

• So each s subshell can contain up to 2 electrons

• first subshell

• Coefficient = period #

p subshell

• Dumbbell shaped

• Each p subshell consists of 3 orbitals

• So each p subshell can contain up to 6 electrons

• second subshell

• Coefficient = period #

d subshell

• double dumbbell-shaped or cloverleaf-shaped (do not have to know this)

• Each d subshell consists of 5 orbitals

• So each d subshell can contain up to 10 electrons

• third subshell

• Coefficient = period # - 1

f subshell

• Each f subshell consists of 7 orbitals

• So each f subshell can contain up to 14 electrons

• fourth subshell

• Coefficient = period # - 2

Orbitals

• Subshells contain orbitals of different shapes

• Each orbital can contain 2 electrons

• The number of orbitals that each subshell are increasing odd numbers

  • the s subshell has 1 orbital, the p subshell has 3 orbitals, the d subshell has 5 orbitals, and the f subshell has 7 orbitals

• How they are filled follows 3 rules

  • Aufbau Principle, Pauli Exclusion Principle, and Hund’s Rule

Nodes

areas in which electrons are not found

Orbital Notation

• A graphical representation of exactly where electrons of an atom are placed.

  • Tells you which shell, subshell and orbital the electrons are in

• Uses arrows to represent electrons and lines to represent orbitals.

  • Up arrow represents clockwise spinning electron and down arrow represents counter-clockwise spinning electron

Aufbau Principle

• Electrons fill orbitals in order of increasing energy (1s, 2s, 2p, 3s, 3p, 4s, 3d, etc.)

  • because that is their most stable, lowest-energy state

Pauli Exclusion Principle

• No two electrons can have the same location and spin

  • Reduces repulsion between the electrons

Hund’s Rule

• All orbitals of a given subshell must contain one electron before pairing begins

  • This is because electrons are negatively charged so they repel each other so as a result they will be a far from each other as possible

Electron Configuration

• Shorter and easier way to represent the placement of electrons in the atom

• No arrows, no orbital lines

  • Does not show spin or the orbitals

• Superscripts are used to represent electrons

• Less specific than orbital notation

Electron Configuration of Anions

• Electrons need to be added to the last subshell or the next subshell if the previous shell fills

• Ex:

  • S : 1s²2s²2p⁶3s²3p⁴

  • S^2- : 1s²2s²2p⁶3s²3p⁶

  • S^3- : 1s²2s²2p⁶3s²3p⁶4s¹

Isoelectronic

2+ atoms/ions that have the same number of electrons and the exact same electron configuration, even though they belong to different elements

Ex:

Ne : 1s²2s²2p⁶

F^- : 1s²2s²2p⁶

O^2- : 1s²2s²2p⁶

Electron Configuration of Cations

• Electrons need to be removed from the valence shell (the subshell with the highest coefficient)

• Ex:

  • Sc : 1s²2s²2p⁶3s²3p⁶4s²3d¹

  • Sc¹⁺ : 1s²2s²2p⁶3s²3p⁶4s¹3d¹

  • Sc²⁺ : 1s²2s²2p⁶3s²3p⁶3d¹

  • Sc³⁺ : 1s²2s²2p⁶3s²3p⁶

Noble Gas Electron Configuration

• Shorter than Electron Configuration

• Put the previous noble gas in brackets and then list the rest of the electron configuration.

• Ex:

  • As: [Ar]4s²3d¹⁰4p³

  • U: [Rn]7s²5f⁴

Chemical Periodicity/Periodic Trends

• Trends that exist in the periodic table

  • Atomic Size/Atomic Radius

  • Ionization Energy

  • Reactivity

  • Electronegativity

Reactivity

• As you go down each group they increase in reactivity

  • because outer valence electrons become easier to lose in larger atoms the outer valence electrons are further away from the positive charge of the nucleus so the experience a weaker nuclear attraction force

Atomic Radius

• Distance from the nucleus to the valence shell

• Vertical: going down each group, atomic radius increases

  • More energy levels/shells so the atom takes up more space

• Horizontal: going left to right in each period the atomic radius decreases

  • The number of protons increases from left to right in each period so the valence electrons experience a greater nuclear force from the protons in the nucleus, pulling in the electron cloud and shrinking the size of the atom

Ionization Energy

• The amount of energy needed to remove an electron. (Energy needed to turn a neutral atom into a cation)

• Vertical: going down each group, ionization energy decreases

  • More energy levels/shells so the electrons in the valence shell are further from the nucleus so they experience a weaker nuclear attraction force from the nucleus so they require less energy to be removed

• Horizontal: going left to right in each period the ionization energy increases

  • The number of protons increases from left to right in each period so the valence electrons experience a greater nuclear force from the protons in the nucleus so they require more energy to be removed