Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Matter, and more! Matter • Anything that has MASS and takes up SPACE How is Matter classified? • 1) Pure Substances • 2) Mixtures 1) Pure Substances • Composition remains the same, does not depend on a sample = Fixed composition • Homogenous—same throughout. • Example: Compound (NaCl) or Element (Fe) 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” – 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 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 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 Law of Conservation of Mass/Matter • Matter cannot be created or destroyed • Total mass is constant in chemical reactions. • Originated with Antoine Lavoister (1700s) – Quantitative mass data of reactants and products in mercury oxide decomposition. 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 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 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 possible charge (-). • Determined ratio between mass and charge of an electron 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 10-31kg • Discovered electron charge – e = -1.602 x 10-19 C 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 Contained protons Negative electrons scattered around the outside 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 as cloud of negative charge 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 ! ! ! 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 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 • Specific for each element • Used for element identification 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 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 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 • 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 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 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 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 8 electrons • n=3 – Holds 18 electrons • n=4 – Holds 32 electrons 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 = 7orbitals, can hold 14 eletrons 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 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 “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 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 value determines size – Charge cloud found near center, likely electron location • p orbital – 3 possible orbital orientations, dumbbell shape – pX, py, pz l Orbital/Subshell 0 s 1 p 2 d 3 f How do you specify orbitals? • 2p • 4f 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 (, ) 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 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 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 electrons • Subshells fill with electrons at different regions within periodic table (s section, p section) Ex. 1 Nitrogen Ex. 2 Mg Practice 1) Boron 2) Copper 3) Sodium The Name’s BOND……. Chemical Bond Diatomic Molecules • Always exist in a chemical bond with another atom, even if the atom is of the same element • H2 , N2 , O2 , F2 , Cl2 , Br2 , I2 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 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 • 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 H2 or Cl2 POLAR COVALENT BONDS when electrons are shared but shared unequally H 2O 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 NaCl, CO METALLIC BOND bond found in metals; holds metal atoms together very strongly 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 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