Download SCIENCE LONG TEST

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