Chemistry Unit 6

States of Matter, Kinetic Molecular Theory (KMT)

Kinetic Theory

  • Kinetic = motion

  • Kinetic Energy = energy an object has because of its motions

  • Kinetic Molecular Theoru = tiny particles in all state of matter are in constant motion (solids vibrate, liquids slide, gas is random). Unless Kelvin is absolute 0, it is moving

Gases

  1. Gases are composed of particles (atoms or molecules)

    a. Empty space between particles

    b. No forces between particles

    c. Gases fill containers regardless of volume

    d. if uncontained, gases diffuse into the atmosphere within gravitational limits

  2. Gas Particles move rapidly,

    a. gases move in straight paths, only change path when collide

  3. Collisions are elastic meaning that kinetic energy is transferred from one particle to another but total KE remains constant.

Gas Pressure

  • Gas pressure is the result of particles colliding with objects

  • A vacuum is an area with no particles

  • Atmospheric pressure is the pressure that air exerts on earth because gravity holds air molecules into the earths atmosphere

  • Pressure increases closer to earth or lower elevation

  • Pressure decreases as elevation increases because there is less air

  • Gas volume decreases with increasing pressure

  • increase heat increases pressure

  • gases are compressable

  • Barrometers measure air pressure

  • Units: 1 atm = 760 mm Hg = 101.3kPa = sea level

  • Permanent gases- are gasses that cannot be turne into liquid with just pressure

Kinetic Energy and Kelvin Temperature

  • When a substance is heated, particles absorb the energy and store it. Stored energy does not raise the temperature.

  • When max energy capacity of particles is stored, remaining energy increases average Kinetic Energy and particles collide and move faster and increases the temperature

  • Kelvin temp is related to kinetic energy

Absolute 0 is 0 Kelvin and no movement, and never has been achived

Liquids

  • particles are in constant motion and slide pass eachother and flow

  • Has some intermolecular forces that hold liquid together

  • Liquids are more dense gases

  • Liquids and Solids are condensed meaning that pressure does not reduce volume

Evaporation/ Vaporization

  • Converstion of a liquid to gas below boiling point

  • When particles at the surface of a liquid have enough energy to break IMFs they escape as a gas and it is cooling process because particles with the highest KE or most temp escape and leave behind the liquids with less KE or cooler.

Boiling Point

  • tempature at which the vapor pressure of the liquid equals the external pressure

  • tempature that substance transitions from liquid to gas

  • tempature of boiling liquid will never exceed boiling point because it would have been a gas

  • normal boiling point = boiling point of liquid at STP

  • Presure changes boiling point. High elevations = lower boiling points because lower pressure

  • Boiling with low elevations or high pressure will need more energy, or higher temp

Solids

  • particles vibrate about fixed points

  • most particles are packed in organized patterns 3-d lattice (cystaline)

  • Dense and condensed (cant be compressed)

  • Applying heat causes more vibration of particles and breaking IMFs. Melting Points

Melting Point/ Freezing Points

  • Melting Point = temp are which solid turns into liquid

  • Freezing point = temp at which liquid turns into solid

  • Freezing point = melting point

  • Ionic solids have high melting point, molecular have low

Allotropes

  • different forms of the same physical element

Amorphous Solids

  • not crystal and has no lattice (pattern)

  • lacks organization

  • an example is glass which is also referred as a super liquid becayse it cools to a rigid state without cystalzying.

  • a super fuid is a liquid that cools to rigid state without crystallizing

  • rubber is also amorphorous

Sublimation

  • vapor pressure of some solids is high enough that they can go from solid to gas directy.

  • opposite of sublimation is depostion

  • ex: dry ice, moth balls

What is Vapor pressure?

Plasma

  • 4th state of matter

  • when gas is heated even more, the molecules will seperate into atoms and then the atoms e- will strip off and positive ions are left with free e-

  • the sun is plasma, and partial plamas include lightning bolts, and neon lights

  • This moves faster than gas

Bose Einstin Condensate

  • 5th state of matter

  • super cold atoms, and particles start to act like waves instead of particles

  • atoms can not be distinguished

  • slow

A Matter of state video

  • Energy is needed to change states of matter. Temperature and pressure

  • temperature and pressure of a gas are related. Increase of tempature increases pressure because it increases the speed of particles and they colide with the walls harder. Converse is true

  • Liquified Natural gas is produced with high pressue and cooling it.

  • Uses of liquid Nitrogen: it can fast freeze things, storage of biological specimens “organs, and refrigerant.

  • Why does perspiration (evaporation) cool a person. Liquid water is heated to gas and leaves cool water because water absorbs the heat.

  • Bromine 3 states gas is orange, liquid is red, yellow solid.

  • Particles in bromine when cooled lose kinetic energy, slow down and condense. Attractive forces increase

  • Crystal shape - organized lattice (pattern), the external shape is the same as the internal molecular shape.

  • amophorous solid - glass, not organied, no lattice (pattern).

A Race for Absolute Zero

  • Tech - MRI, Refiferator, Conditions

  • Permanent Gases - oxygen, hydrogen, and nitrogen

    • they cant liquidity with just pressure

  • superconductivity - all resistance to he flow of electricity stops at cold temperatures

  • we will probably not reach 0k soon

Lewis Dot Structures

  • Covalent Bonds - 2 or more electrons are shared by 2 atoms to complete their octet

    • double bonds

    • triple bonds

    • Occurs between 2 non metals

  • Ionic Bonds - the metal will give up electrons to the non metal and the charges create a bond

    • more temp needed to break the bonds “stronger”

  • Steps

    • Draw the skeletal structure “most electronegative in the center”

    • Count the e-, consider the charges too

    • complete octect for all except hydrogen

    • too many electrons → use double or triple bonds

    • consider # of bonds

  • resonance structure - one multiple lewis structures can represent a molecule

Exceptions

  • Incomplete Octet:

    • Be only needs 4e-, or 2 pairs

    • B only needs 3 pairs because has 3 valence 3e-

  • odd Electron Molecules

  • Expanded Octet - can occur with central atoms with principal quantum number more than 2. third-row (or higher)

    • place extra pairs around the atom

    • use empty valence d orbitals

  • CNOF always follow, second row never exceed

Shapes or Electronic Geometry VSEPR

  • VSEPR - theory that predicts the shape based on electron pairs repelling

  • Valence

  • Shell

  • Electron

  • Pair

  • Repulsion

VSEPR Theory - Postulates, Limitations, Predicting Shapes

Polarity

  • Electronegativity (EN) is an atom's tendency to attract electrons in chemical bonds.

    EN increases to the right and up on the periodic table, excluding the noble gases.

    • side to side matters more than vertical

Bond Polarity

  • When two nonmetal atoms bond, they share electrons. They may or may not share electrons evenly.

  • If Same atom, then it is non polar because the non has more favor for electrons

    • highest electron density between atoms

  • if different atoms, one will be more electronegative and will pull electrons more thus its called a polar bond

    • creates partial charges

    • e- density is drawn towards the side of the more electronegative atom

  • Dipole Moments - is a measure of the polarity of bonds and is represented by an arrow which points to the more EN atom

  • Polar does not mean charges

Polarity in molecules of 3 or more Atoms

  • Polar if the central atom has lone pairs or the outer atoms are not all the same

  • nonpolar if the central atoms have no lone pairs and the outer atoms are identical

Properties

  • like disolve like

  • Polar molecules have a higher melting and boiling point than non polar because polar can have IMF - still far below than ionic

Intermolecular Forces - force between 2 molecules but not bonding

  • dipole dipole interaction - is when the partial charges attract → higher melting and boiling point

  • hydrogen bonding - is a stronger dipole interaction

    • occurs in molecules that have H bonded to ONF

    • hydrogen is small so it allows for the molecules to be close → stronger

  • london dispersion forces - the temporary attraction between dipoles

    • can happen for anything

    • more e- → stronger force

Radioactivity

General

  • Unstable isotopes try to gain stability by making changes

    • these changes release lots of energy

  • reactions are not affected by temperature, pressure, catalysis, or compound

  • cannot be slowed of speeded up

  • Becqueral, Marie and Pierre Curie

Terms

  • Radioisotope: radioactive isotope has too many neutrons

  • Radioactive decay: the process by which an unstable nucleus loses energy by emitting radiation “spontanous” → to a isotope of a different element

Types of Radiation

  • Alpha (α) 24H – a positively charged helium isotope

    • weakest

    • The mass number goes down by 4, the atomic number goes down by 2

  • Beta (β) – an electron – formed from decomposition of a neutron

    • medium strength

    • the mass number is unchanged and the mass number goes up by one

    • (one less neutron, one more proton)

  • Gamma (γ) – pure energy; called a ray rather than a particle

    • strongest, most dangerous and can affect DNA

Half Life - the time it takes for ½ of a sample to decompose

  • the rate of nuclear transformation depends on the reactant concentration

  • Mass and T table

    • start with the most mass and 0 time

    • divide the mass by 2 til the final value is reached

    • divide the time divided by the time

Transmutation: conversion of atoms of one element into another element

  • may be natural or manmade

  • all isotopes above #82 is radioactive

  • Transuranium element: man made and over #92

Nuclear Fission -

  • “spliting of atoms”

    • usually in large atoms that are not as stable

  • releases lots of energy

Ionizing Radiation - radiation strong enough to knock off electrons of the atoms that it strikes

Nuclear reactors use controlled fission to produce energy as heat, and this makes steam that turns turbines and generates electricity

  • Nuclear reactions are powered by bombarding neutrons and a chain reaction - the atoms will separate and produce neutrons which will bombard other atoms

  • Controlling Fission: to prevent too much heat or explosions → meltdowns

    • Neutron Moderation: slows neutrons with water and carbon

    • Neutron Absorbtion: traps neutrons in control rods

    • Cherynoble - lack of control

  • makes lots of energy

Nuclear Fusion

  • its the opposite of fission

  • Fusion - small nuclei combine “hydrogen combine" → helium”

  • This occurs in the sun and stars which created lots of energy

  • Exessive heat and hard to contain

Detecting Radiation

  • Can detect results of Radiation

  • Geiger Counter: tube with neutral gas

    • best for beta “e-”

    • it interacts with the gas inside, causing ionization.

  • Scintillation Counter: flasg of light occur when ions strike the phosphor

    • works for all radiation

    • best for alpha

  • Film Badge: is a photographic film encased in holder - used by workers

Research and Medicine

  • Neutron Activation: determine the age of a sample

  • Radioisotope Tracers: used to follow chemicals throughout a bioloifcal process

  • Cancer Treatment: since cancer cells divide more rapidly, they are susceptible to radiation

    • telethearapy

    • implantation

    • pharmacueticals