Matter and more Matter Anything that has MASS

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Matter, and more!

Matter, and more!

Matter • Anything that has MASS and takes up SPACE

Matter • Anything that has MASS and takes up SPACE

How is Matter classified? • 1) Pure Substances • 2) Mixtures

How is Matter classified? • 1) Pure Substances • 2) Mixtures

1) Pure Substances • Composition remains the same, does not depend on a sample

1) Pure Substances • Composition remains the same, does not depend on a sample = Fixed composition • Homogenous—same throughout. • Example: Compound (Na. Cl) or Element (Fe)

2) Mixtures • 2+ substances combined together • Substances do not change their properties

2) Mixtures • 2+ substances combined together • Substances do not change their properties or name. • Able to be separated, not chemically combined. • Possess a combination of properties based on the substances present.

Types of Mixtures 1) Homogenous – Uniform composition – Also known as “true solutions”

Types of Mixtures 1) Homogenous – Uniform composition – Also known as “true solutions” – Ex. Salt-water 2) Heterogeneous – No uniform composition – Can easily see the different components of the mixture – Ex. Italian dressing

Properties of Matter 1) Chemical – Ability to go through changes resulting in a

Properties of Matter 1) Chemical – Ability to go through changes resulting in a different substance – The substance is no longer the same, different identity – Ex. Burning 2) Physical – Observed or measured property – Substance identity is not changed – Ex. melting point, boiling point, density

Practice

Practice

Extensive vs. Intensive Properties • Extensive Property – Dependent on the quantity of matter

Extensive vs. Intensive Properties • Extensive Property – Dependent on the quantity of matter – Ex. Mass • Intensive Property – Does NOT depend on matter quantity – Ex. Density

The Atom

The Atom

Law of Conservation of Mass/Matter • Matter cannot be created or destroyed • Total

Law of Conservation of Mass/Matter • Matter cannot be created or destroyed • Total mass is constant in chemical reactions. • Originated with Antoine Lavoister (1700 s) – Quantitative mass data of reactants and products in mercury oxide decomposition.

Terminology • Element– basic unit of a substance, contain only ONE type of atom,

Terminology • Element– basic unit of a substance, contain only ONE type of atom, represented by symbol. Example: Ag, only contains Ag atoms. • Atom—smallest particle of an element that still contains element properties. – Example: One atom of Au, cannot have a smaller particle of gold and still be gold.

Compound vs. Molecule • Compounds: – more than one element – elements combined in

Compound vs. Molecule • Compounds: – more than one element – elements combined in definite proportions • Molecule: – Smallest unit of a compound that still retains the properties of the compound.

Dalton Atomic Theory (cont. ) • All matter made of atoms • Atoms of

Dalton Atomic Theory (cont. ) • All matter made of atoms • Atoms of an element have the same size, mass, etc. • Different atoms have various sizes, mass, etc. • Atoms cannot be divided, destroyed, or created. • Atoms rearrange in chemical reactions.

John Thomson • 1897 • Cathode-Ray experiments. • Discovered the electron particle and its

John Thomson • 1897 • Cathode-Ray experiments. • Discovered the electron particle and its possible charge (-). • Determined ratio between mass and charge of an electron

Robert Millikan • Oil drop experiments. • 1909, American • Found the mass of

Robert Millikan • Oil drop experiments. • 1909, American • Found the mass of an electron (VERY small) with Thompson’s data • Currently, mass of electron = 9. 109 x 10 -31 kg • Discovered electron charge – e = -1. 602 x 10 -19 C

Early Models of the Atom Thompson • Must be a balance between negative and

Early Models of the Atom Thompson • Must be a balance between negative and positive charges • “Raisin-Pudding” model • Uniform distribution of positive charge – Positive cloud with stationary electrons

Early Models of the Atom Rutherford • • Mostly empty space Small, positive nucleus

Early Models of the Atom Rutherford • • Mostly empty space Small, positive nucleus Contained protons Negative electrons scattered around the outside

Early Models of the Atom Bohr • 1913—hydrogen atom structure • Physics + quantum

Early Models of the Atom Bohr • 1913—hydrogen atom structure • Physics + quantum theory • Electrons move in definite orbits around the positively charged nucleus—planetary model • Does not apply as atoms increase in electron number

Heisenberg’s Uncertainty Principle • Electron’s location and direction cannot be known simultaneously • Electron

Heisenberg’s Uncertainty Principle • Electron’s location and direction cannot be known simultaneously • Electron as cloud of negative charge

Modern Model of the Atom The electron cloud • Sometimes called the wave model

Modern Model of the Atom The electron cloud • Sometimes called the wave model • Electron as cloud of negative charge • Spherical cloud of varying density • Varying density shows where an electron is more or less likely to be

How did we discover electron arrangement in an atom? ELECTROMAGNETIC RADIATION ! ! !

How did we discover electron arrangement in an atom? ELECTROMAGNETIC RADIATION ! ! !

Electromagnetic Waves • Produced from electric charge movement • Changes within electric and magnetic

Electromagnetic Waves • Produced from electric charge movement • Changes within electric and magnetic fields carried over a distance • No medium needed

Types of Spectra • What is a spectra? – Spectrum– white light/radiation split into

Types of Spectra • What is a spectra? – Spectrum– white light/radiation split into different wavelengths and frequencies by a prism • Continuous spectrum – No breaks in spectrum – Colors together • Line spectrum – Line pattern emitted by light from excited atoms of a particular element – Aided in determining atomic structure

Line Spectrum • Pattern emitted by light from excited atoms of an element •

Line Spectrum • Pattern emitted by light from excited atoms of an element • Specific for each element • Used for element identification

THEREFORE ……… • Light is in the form of electromagnetic waves • Photons can

THEREFORE ……… • Light is in the form of electromagnetic waves • Photons can resemble particles • Gave raise to the possibility of thinking about wave AND particle qualities of subatomic particles (electron)

Atomic Structure • Nucleus – Protons – Neutrons • Electrons

Atomic Structure • Nucleus – Protons – Neutrons • Electrons

Describing Atoms • Atomic Number = number of protons • In a neutral atom,

Describing Atoms • Atomic Number = number of protons • In a neutral atom, the # of protons = the # of electrons • Atomic Mass= the number of protons + the number of neutrons

Isotopes • The number of protons for a given atom never changes. • The

Isotopes • The number of protons for a given atom never changes. • The number of neutrons can change. • Two atoms with different numbers of neutrons are called isotopes • Isotopes have the same atomic # • Isotopes have different atomic Mass #’s

Ions • An atom that carries an electrical charge is called an ion •

Ions • An atom that carries an electrical charge is called an ion • If the atom loses electrons, the atom becomes positively charged. • If the atom gains electrons, the atom becomes negatively charged

PEN Method for-- • • F Li +1 Cl -1 Be

PEN Method for-- • • F Li +1 Cl -1 Be

Light • Composed of small energy packets (photons) • Quantum = minimum amount of

Light • Composed of small energy packets (photons) • Quantum = minimum amount of energy lost/gained by atom • Atoms can absorb or give off this energy

Energy States in an Atom • Atoms can gain or loss energy. • Specific

Energy States in an Atom • Atoms can gain or loss energy. • Specific energy states within an atom. – Can be counted – Ground State = lowest energy state – Excited State = higher energy level than ground, gained energy

So, where does the Bohr Model fit in? • Electrons orbit around the nucleus

So, where does the Bohr Model fit in? • Electrons orbit around the nucleus at different energy levels/orbits. • Electron’s energy level = orbit level where electron is located. • Light absorption = electron moves from a state of low energy to high energy. “becomes excited” • Light Emitted = electron falls from an “excited” state of energy to a lower energy level.

Main Energy Levels/Electron Shell • n=1 – Holds 2 electrons • n=2 – Holds

Main Energy Levels/Electron Shell • n=1 – Holds 2 electrons • n=2 – Holds 8 electrons • n=3 – Holds 18 electrons • n=4 – Holds 32 electrons

Energy sublevels • Within the main energy level. • • S = 1 orbital,

Energy sublevels • Within the main energy level. • • S = 1 orbital, can hold 2 electrons p = 3 orbitals, can hold 6 electrons d = 5 orbitals, can hold 10 electrons f = 7 orbitals, can hold 14 eletrons

Quantum Theory • Treats electron’s location as wave property • Defined by quantum numbers

Quantum Theory • Treats electron’s location as wave property • Defined by quantum numbers • Orbitals have different energies • Quantum numbers – Provide information about size, shape, and orientation of atomic orbitals – Define atomic orbitals from general to specific

Quantum Mechanical Model • Opposite charges attract, electrons are attracted to the nucleus of

Quantum Mechanical Model • Opposite charges attract, electrons are attracted to the nucleus of an atom – Takes a LOT of energy to keep electrons away from the nucleus. • Electrons are found at differing lengths from the nucleus and can only be present in certain locations

Principal Quantum Number (n) • Determines orbital size and electron energy • Same as

Principal Quantum Number (n) • Determines orbital size and electron energy • Same as “n” value/orbital in Bohr model • Positive whole number, NOT 0 • Shells – orbitals with same value • n = 1, 2, 3, 4, etc.

Orbital Angular Momentum Quantum Number (l) • Defines orbital shape for a particular region

Orbital Angular Momentum Quantum Number (l) • Defines orbital shape for a particular region of atom • Think of as “subshell” • Energy sublevels—within the main energy level – s = 1 orbital, can hold 2 electrons – p = 3 orbitals, can hold 6 electrons – d = 5 orbitals, can hold 10 electrons – f = 7 orbitals, can hold 14 electrons –

Orbital Shapes • s orbital – 1 possible orbital orientation, spherical shape – n

Orbital Shapes • s orbital – 1 possible orbital orientation, spherical shape – n value determines size – Charge cloud found near center, likely electron location • p orbital – 3 possible orbital orientations, dumbbell shape – p X, p y , p z

l Orbital/Subshell 0 s 1 p 2 d 3 f

l Orbital/Subshell 0 s 1 p 2 d 3 f

How do you specify orbitals? • 2 p • 4 f

How do you specify orbitals? • 2 p • 4 f

Electron Configuration • Shorthand method for representing electrons’ distribution in orbitals within subshells •

Electron Configuration • Shorthand method for representing electrons’ distribution in orbitals within subshells • All orbitals have the same energy level—digenerate • Orbitals– mathematical expressions of probability of electron’s location • Electrons occupy orbitals in a way that gives LOWEST energy state

Orbital Diagrams • Visual representation of electron configuration • Represents electrons’ spins ( ,

Orbital Diagrams • Visual representation of electron configuration • Represents electrons’ spins ( , )

Aufbau Principle • Electrons occupy the LOWEST energy orbital available • Lazy Hogs !

Aufbau Principle • Electrons occupy the LOWEST energy orbital available • Lazy Hogs !

Hund’s Rules cont. 1) One electron MUST occupy each orbital BEFORE electrons are paired

Hund’s Rules cont. 1) One electron MUST occupy each orbital BEFORE electrons are paired in the same orbital. 2) Electrons added to subshell with the same spin (+1/2, -1/2) so each orbital has one electron.

Pauli Exclusion Principle • Only 2 electrons occupy each orbital • Electron spins MUST

Pauli Exclusion Principle • Only 2 electrons occupy each orbital • Electron spins MUST be opposite/paired when 2 electrons occupy the same orbital – +1/2, -1/2

Using the periodic table- • Period numbers = principal quantum number of valence shell

Using the periodic table- • Period numbers = principal quantum number of valence shell electrons • Subshells fill with electrons at different regions within periodic table (s section, p section)

Ex. 1 Nitrogen

Ex. 1 Nitrogen

Ex. 2 Mg

Ex. 2 Mg

Practice 1) Boron 2) Copper 3) Sodium

Practice 1) Boron 2) Copper 3) Sodium

The Name’s BOND……. Chemical Bond

The Name’s BOND……. Chemical Bond

Diatomic Molecules • Always exist in a chemical bond with another atom, even if

Diatomic Molecules • Always exist in a chemical bond with another atom, even if the atom is of the same element • H 2 , N 2 , O 2 , F 2 , Cl 2 , Br 2 , I 2

Chemical Bond • Involves 2 atoms • Mutual, electrical attraction between nuclei and valence

Chemical Bond • Involves 2 atoms • Mutual, electrical attraction between nuclei and valence electrons of 2 atoms

IONIC BOND bond formed between two ions by the transfer of electrons. ELECTRON STEALING

IONIC BOND bond formed between two ions by the transfer of electrons. ELECTRON STEALING !

Ionic Bonds • Between atoms of metals and nonmetals with very different electronegativity –

Ionic Bonds • Between atoms of metals and nonmetals with very different electronegativity – Huge difference in electronegativity – Electronegativity value > 1. 7 • Bond formed by transfer of electrons • Produce charged ions all • Form between elements located on opposite sides of the periodic table

COVALENT BOND bond formed by the sharing of electrons Electron sharing!

COVALENT BOND bond formed by the sharing of electrons Electron sharing!

Covalent Bond • Between nonmetallic elements of similar electronegativity. – Small electronegativity difference <1.

Covalent Bond • Between nonmetallic elements of similar electronegativity. – Small electronegativity difference <1. 7 • Formed by sharing electron pairs, • Formed between elements on the same side of the periodic table.

NONPOLAR COVALENT BONDS when electrons are shared equally, small electronegativity difference H 2 or

NONPOLAR COVALENT BONDS when electrons are shared equally, small electronegativity difference H 2 or Cl 2

POLAR COVALENT BONDS when electrons are shared but shared unequally H 2 O

POLAR COVALENT BONDS when electrons are shared but shared unequally H 2 O

Polar Covalent Bonds--More • One atom “keeps” electrons closer to it • Electrons tend

Polar Covalent Bonds--More • One atom “keeps” electrons closer to it • Electrons tend to reside around one atom more than the other atom • Electrons still remain distributed between the 2 atoms--unequal

Identify Bond type, Draw Lewis Structure to show bonding Na. Cl, CO

Identify Bond type, Draw Lewis Structure to show bonding Na. Cl, CO

METALLIC BOND bond found in metals; holds metal atoms together very strongly

METALLIC BOND bond found in metals; holds metal atoms together very strongly

Electrons • Localized: – Electrons hang out in a local area – Restricted to

Electrons • Localized: – Electrons hang out in a local area – Restricted to an atom/ion or shared between atoms • Delocalized: – Electrons not attached to a particular area – Electrons are able to travel and move among atoms

Metallic Bonds • Formed between atoms of metallic elements – Between atoms on left

Metallic Bonds • Formed between atoms of metallic elements – Between atoms on left side of periodic table – Electronegativities are approximately equal • Electron cloud around atoms, delocalized electrons – Valence electrons shared among multiple atoms, “neighbors” • Electrons move through the whole metal, can jump between energy levels to create conduction bands.

Metallic Bond, A Sea of Electrons

Metallic Bond, A Sea of Electrons