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SCIENCE LONG TEST 4TH GRADING I. Concepts: Matter A. Matter – is composed of atoms B. Atoms – come from the greek word atomos which means indivisible C. 1. 2. 3. 4. Early ideas about the atom Atom – comes from the greek word aromos which means indivisible Democritus and Leucippus – first to propose that matter is made up of atoms Democritus’ hypotheses atoms were small, hard particles made of the same material but of different shapes and sizes there were an infinite number of these atoms and they were constantly in motion atoms had the ability to combine with other atoms atoms could no longer be divided into smaller particles The early ideas about the atom were quite close to what we know today D. 1. 2. 3. 3 subatomic particles/symbols/charges/discovered by protons – positively charged particles electrons – negatively charged particles neutrons – have the same mass as the protons but have no electric charge Particles a. nucleus b. protons c. neutron d. electron Symbol p n e/e Charge Positive Positive No charge Negative Discovered by Sir Ernest Rutherford Eugene Goldstein James Chadwick Joseph John Thompson E. 1. 2. 3. 4. 5. 6. 7. 8. Elements, Molecules, and Compounds 92 different kinds of atoms make up all naturally occurring substances Less than a hundred kinds of atoms combine to form all the things you will ever get to know 21 atoms have been artificially produced in laboratories In total, 113 different kinds of atoms are there in our universe Elements – substances that are made up of only one kind of atom Compound – substances made up of more than one type of atom Molecule – smallest piece of a compound; More generally, the molecule is the smallest piece of a pure substance – substances that always have the same component particles in the same proportion. It is made up of two or more atoms 9. Most substances found in nature occur as molecules. Very few exist as isolated atoms F. 1. 2. 3. 4. 5. Properties of Atoms Atoms are incredibly small Everything around us is made up of things too small to be seen may be hard to understand and believe Atoms are eternal Atoms are continuously being recycled Cell – basic building block of living things G. Different Models of the Atom Dalton’s Model =>Smallest indivisible particle of mater Thomson’s Model =>Made up of even smaller pieces =>Consisted of a positive material scattered with negatively charged particles Rutherford’s Model => Consisted of a positively charged center which also contained most of the atom’s mass Bohr’s model =>Placed the electrons in orbit at specific distances around the nucleus. This resembled the way in which planets revolve around the sun Quantum Mechanical Model: present model of an atom => The location of the electrons around the atom cannot be precisely determined. The region where the electron can probably be found is known as the electron cloud H. 1. 2. 3. 4. 5. Terminologies atomic mass number – an arbitrary unit that will express the mass of an atom atomic number – number of protons; atomic number mass number – sum of protons and neutrons (atomic mass); number of nucleons isotopes – atoms with the same elements with different atomic numbers of neutrons energy level – represent a particular location within the electron cloud where an electron within an energy level 6. energy sublevel – located in the main energy level 7. electron orbital – region of space around the nucleus where electrons are most likely to be found 8. neutral atom – equal number of positive and negative charges I. Parts of an Element A B D F C E G H A. B. C. D. E. F. G. H. I. I IJ mass no. atomic no. electron configuration notation main energy level sublevel Number of electrons Last Main Energy Level – consider the biggest period Valence Electrons – group no. Group/Family J. Getting the values of an element Atomic number Atomic mass 1. 2. 3. 4. Protons Electrons Neutrons PROTONS: SAME AS ATOMIC NUMBER AND ELECTRONS NEUTRONS: AM-AN ELECTRONS: SAME AS PROTONS ATOMIC MASS: PROTONS+NEUTRONS LMEL – biggest period Valence Electrons – group no. K. Main Energy Levels/Sublevels/Orbitals/Electron Capacity Main Energy Levels Sublevels Orbitals => 1 S S=1 2 S&P S=1 P=3 3 S, P & D S=1 P=3 D=5 4 S, P, D, & F S=1 P=3 D=5 F=7 Group-Family – same as valence (s/p=A) (d/f=B) Electron Capacity (x2) Max of 2 electrons Max of 8 electrons Max of 18 electrons Max of 32 electrons L. Periodic Table of the Elements 1. Period/Series (rows of elements) – horizontal arrangement of elements in the periodic table [main energy levels] 2. Groups/Families (columns of elements) – vertical arrangement of elements in the periodic table o 1. 2. 3. 4. 5. 6. 7. S P D F 7. 8. Structure of the Periodic Table The periods or series refer to the main energy levels of an atom (1-7 periods) Atomic numbers of elements increase from L-R of the P.T. Metals are located on the left side of the zig-zag line Non-metals are located on the upper right side of the zig-zag line The metalloids bonder the zig-zag line Groups or Families A Families [IA to VIIIA] B Families [IB to VIIIB] Inert [VIIIA] Sublevels 2e Sharp lines 3 orbitals, max of 6e Principal lines 5 orbitals, max of 10e Diffused lines 7 orbitals, max of 14e Fundamental lines Inner – lanthanoids [period6] and actinoids [period7] (f blocks) Outer – s, p, d blocks M. Principles 1. Aufbau’s Building-up principle Electrons closer to nucleus – lesser amount of energy Electrons farther to nucleus – more amount of energy 2. Pauli’s Exclusion Principle Repulsion – 2 electrons of the same charge, there would be a tendency of repulsion 2 electrons spin in opposite directions thus developing opposite magnetic directions 3. Hund’s rule of maximum multiplicity Each electron occupies an orbital one at a time before pairing takes place N. Notations Electron Configuration Notation 37 Rb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 5s2 4d5 When crossing to d block (n-1) deduct 1 to the period When crossing to f block (n-2) deduct 2 to the period Finding the ve: When there are two highest periods, simply add the number of electrons Orbital notation – subscript is not included; ECN is needed 1s2 2p6 3d10 6f14 Electron dot notation – Lewis Symbol; Gilbert Lewis (proponent) 3 6 7 4 1 2 8 5 II. Chemical Bonding A. Introduction of Terms 1. Valence Electron - total no. of electrons in the last main energy level - no. of electrons available for bonding 2. Valence Shell - energy level where the valence electrons are located 3. Chemical Bonding - way where atoms give/share their valence electrons to form compounds B. Types of Chemical Bonding Goal: To make the element a noble/inert gas 1. Ionic bonding – gain or lose electrons; metals and non metals Octet Rule – standard in which elements follow in bonding Noble Gases – elements with completely filled orbital Electron affinity – attraction of electrons Oxidation – process when atoms lose electrons Reduction – process when atoms gain electrons Gilbert Lewis a physicist gave the Lewis dot symbol shown on the right Ionic substances have high melting and boiling points Ionic Compounds – formed element Oxidation/Reduction Diagram 2. Covalent bonding – sharing of electrons; both non metals 1. Hydrogen – prone to covalent bonding 2. Gilbert Lewis reasoned out that atoms could also obtain noble gas electron configurations by sharing their electrons 3. No one gains/loses electrons 4. The bonding pair of electrons is usually shown in a line and the unshared electrons as dots 5. Electronegativity - starting from hydrogen going to the right (increases) and going down (decreases) - relative ability of two atoms to attract bonding electrons to themselves Types of Covalent Bonding a. Non polar covalent bonding – equal sharing of electrons b. Polar covalent bonding – unequal sharing of electrons; one element attracts more orbital than the other H Cl Dipole/bond polarity – equal distribution of charges Di = two Pole = positive/negative c. Multiple Covalent Bonds H O H (2 parts H) (1 part oxygen) H O H Representing a Covalent Bond – Venn diagram C. Principles on Bonding [applications] Octet Subrules a. Metals lose 1-3 ve to form a cation with the structure of a previous noble gas => Cation – similarity of a metal losing electrons to the stability of a noble gas b. Non metals gain 1-3 ve to form an anion with the same structure of the next noble gas => Anion – similarity of a non metal gaining electrons to the stability of a noble gas Noble Gases – 8 valence electrons List of noble gases in order a. Helium (2 electrons; completely filled orbital), b. Krypton c. Radon d. Xenon e. Argon f. Neon Illustrating the bond use “line” illustration when ionic bond use “venn diagram” illustration when covalent bond III. Oxidation # and Chemical Names A. Oxidation Number Getting the Oxidation Number – Family A Representative Family Oxidation Number Lose/Gain IA +1 Lose IIA +2 Lose IIIA +3 Lose VA -3 Gain VIA -2 Gain VIIA -1 Gain *Group IV-A is capable of forming covalent electrons *Group VIII-A is already a noble gas/inert gas Getting the Oxidation Number – computation Element Valence Computation Electron 2 20+p + 20-e = 0 20Ca 20+p + 18-e = +2 lose 6 8+p + 8-e = 0 8O gain B. 8+p + 10-e = -2 Oxidation Number + Ca 2 O-2 Chemical Formula Writing Guidelines: 1. Determine the numerical charge of the ions in the binary compound 2. Cross over the numerical charge of each ion. Write them as subscripts *Element A will exchange its oxidation # with Element B (vice versa) *The Numerical Subscript 1 is not written *Express the ions in their lowest whole number ratio *Ionic Bonds – Left (+) Right (-) *Covalent Bonds – Left (-)[+] Right (-) *In Covalent Bonds. Elements in the left will be changed to positive temporarily C. Chemical Name Writing Binary ionic Name the cation by its atomic name; followed by the stem of the anion’s name in which –ide ending is attached Binary covalent Use the Greek prefixes for the + elements. Then add the Greek prefixes followed by the stems of the – elements with the –ide endings Greek Prefixes Mono Di Tri Tetra Penta Hexa Hepta Octa Nona Deca No. of Atoms 1 2 3 4 5 6 7 8 9 10 Mono is not added to the 1st element Names of Elements: Family S/P Symbol Name of New Name Element H Hydrogen Hydride He Helium metal Li Lithium metal Be Beryllium metal B Boron metal C Carbon Carbide N Nitrogen Nitride O Oxygen Oxide F Fluorine Fluoride Ne Neon Neonide Na Sodium metal Mg Magnesium metal Al Aluminum metal Si Silicon Silide P Phosphorus Phosphide S Sulfur Sulfide Cl Chlorine Chloride Ar Argon Argide K Potassium Metal Ca Calcium Calcide Ga Gallium Metal Ge Germanium Metal As Arsenic Arsenide Se Selenium Selenide Br Bromine Bromide Kr Krypton Kryptide Rb Rubidium Rubide Sr Strontium Strontide In Indium Metal Sn Tin Metal Sb Antimony Antide Te Tellurium Metal I Iodine Iodide Cs Cesium Ceside D. Metals with variable oxidation status Metals Stock Method Classical Method Co+1 Cobalt (I) Cobaltous Co+2 Cobalt (II) Cobaltic Cu+1 Copper (I) Cuprous Cu+2 Copper (II) Cupric Cr+2 Chromium (II) Chromous Cr+3 Chromium (III) Chromic Au+1 Gold (I) Aurous Au+3 Gold (III) Auric Fe+2 Iron (II) Ferrous Fe+3 Iron (III) Ferric Pb+2 Lead (II) Plumbous Pb+3 Lead (IV) Plumbic Sn+2 Tin (II) Stannous Sn+4 Tin (IV) Stannic Lower Charge – ous Greater Charge – ic Polyatomic ions – units and bound groups of atoms that behave as one and carry a charge E. Ion NH OH ClO ClO NO NO HCO CO SO SO +1 4 -1 Polyatomic name Ammonium Hydroxide -1 Chlorate 3 -1 Chlorite 2 -1 Nitrate 3 -1 Nitride 2 -1 Bicarbonate 3 -2 Carbonate 3 -2 Sulfate 4 -2 Sulfite 3 Lower – ide ending Greater – ate ending *refer to the lower charge to determine if it is lower/greater Criss Crossing: 1. when a polyatomic ion gains a subscript of 2 and above, the ion is enclosed in a parenthesis and then write the subscript 2. if it gains a subscript of 1, no parenthesis needed and the subscript [ 1] is not written Cancelling: 1. Ba2 (NO2)4 The subscript 2 and 4 is cancelled; therefore it will have the final answer: Ba (NO2)2 Naming: 1. Format: Monoatomic Name + Polyatomic Name Example: Ba (NO2)2 Barium Nitrite