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slide1

PHYSICAL SCIENCE

SLT STUDY GUIDE

Chemistry and Physics

2012-2013

how to convert between american and si
How to convert between American and SI
  • Dimensional Analysis (also called Factor-Label Method or the Unit Factor Method) is a problem-solving method that uses the fact that any number or expression can be multiplied by one without changing its value. It is a useful technique.
  • Unit factors may be made from any two terms that describe the same or equivalent "amounts" of what we are interested in. For example, we know that

1 inch = 2.54 centimeters

  • We can make two unit factors, each equaling 1 over 1 (or 1) from this information:
slide3

Now, we can solve some problems. Set up each problem by writing down what you need to find with a question mark. Then set it equal to the information that you are given. The problem is solved by multiplying the given data and its units by the appropriate unit factors so that only the desired units are present at the end.

  • For example: You want meters, and are given the data in inches. So –

Given Unit (inches) x Wanted Unit (meters)

Equal Amount in Given

Units (# inches in a meter)

The inches cancel-out leaving your answer in meters

slide4

Example:

    • Note the fraction you multiply by is equal to

one over one. (1 inch = 2.54 cm)

= You use an one/one fraction each time regardless of what the conversion is!

= This is why learning the basic equivalences will make life easier

slide5

More examples: How many seconds in two years?

  • What is the density (D) of mercury (13.6 g/cm3) in units of kg/m3?
atomic models
ATOMIC MODELS

Protons (+ charge) and Neutrons (0 charge) make up the nucleus; Electrons (- charge) around nucleus and held by electromagnetic force. p+ and n made up of quarks.

Plum Pudding Model Electron Cloud Model

isotopes
ISOTOPES
  • All atoms of an element have the SAME number of protons (p+)
  • The p+ number is the atomic number (Z)
    • This is a constant – it stays the same for that element’s atoms
    • For example: All Sodium (Na) atoms have 11 p+
    • If an atom loses a proton, it becomes a different element
      • If Na loses 1 p+, then it has become Neon (Ne)
z atomic number p
Z = atomic number = p+
  • The number of protons identifies the atom and which element it is
  • In a stable atom:
    • # p+ = # n0 = # e-
    • Thus, Na in its stable form has 11 p+; 11 n0; and 11 e-
    • If it has an unequal number of p+ and n0, then it is called an ISOTOPE
slide11
IONS
  • Ions are when an atom has an unequal number of p+ and e-
  • Metals form (+) ions and nonmetals (-) ions
  • Remember – a stable atom has a neutral overall charge due its equal number of p+ and e-
  • When an atom loses or gains an e-, its charge changes accordingly
    • Loss of e- means a + charge; gaining an e- means a – charge for the atom
losing or gaining e
Losing or Gaining e- . . . . .
  • If an atom loses an e-, then it has more p+ than e- and it will have an overall positive charge
  • Different elements’ atoms can lose 1, 2, 3, or even 4 electrons depending on various factors
  • If an atom has LOST e-, then it is called a CATION or a positive ion
    • A Cation would be written as Al+ (the one being understood) or Al+3
slide13

Atoms can also gain electrons

  • If an atom gains electrons (from 1 up to 4), then it will have more e- than p+ and will end up having an overall negative charge
  • A negatively charged ion is called an ANION
  • (A positively charged ion is called a CATION)
  • The NOBLE GASES will not form ions and thus will not bond
  • The Transition Metals can form various numbers of positive ions – got to learn these!
  • The losing or gaining of electrons determines what type of bonds the atoms will form, and which atoms will bond to others
using the periodic table
Using the Periodic Table
  • Elements in the Main Groups (A), form fairly consistent ions
  • Group IA will form +1 ions; Group 2A form up to +2; Group 3A form up to +3 ions
  • Group 4A will form either up to -4 or +4 ions
  • Group 5A will form up to -3 ions; Group 6A up to -2; Group 7A form -1; and Group 8A will not form ions at all
  • Those elements in the B group (transition metals) vary in their + charges meaning they can form different ions
ions and isotopes in review
Ions and Isotopes in Review
  • Stable atom: #p+ = #n0 = #e-
  • Atomic Mass: #n0 = # p+
  • If charge is 0, then #p+ = #e-
  • If charge is positive, then #p+ > #e- Cation
  • If charge is negative, then #p+ < #e- Anion
  • Mass measured in AMU (Atomic Mass Units) based on the C-12 atom
examples
Examples:
  • Li-1 has gained an electron, meaning there is one more negative charge than positive ones
    • It has 3 p+ and 4 e-
  • Li+1 has lost an electron, meaning there is one more positive charge than negative ones
    • It has 3 p+ and 2 e-
      • REMEMBER: The # of p+ DOES NOT CHANGE
      • Only the number of n0 (isotope) and e- (ion) change
slide19

Cf-3 has an atomic number of 98

    • This means it has 98 p+
    • Its atomic mass is 216
    • It has 118 n0, (216 – 98), making it an ion and an isotope!
    • Since it has a -3 charge, the number of e- will be 101; (98 + 3)
    • Zn+1 has 30 p+ and n0; but due to the +1 charge, it has only 29 e-
on the periodic table
On the Periodic Table:

The top number is Z, the Atomic Number or number of p+

The Element’s Symbol

The element average atomic weight set by isotopes and abundances

counting atoms in a molecule
Counting Atoms in a Molecule

In the example, NH3, the subscript 3 only applies to the hydrogen.

  • Therefore: there is 1 N and 3 H in ammonia

In the example, 3Ca3(PO4)2, the number of atoms changes due to the Coefficientis always in front of the whole molecule!!

-The subscript 2, multiplies the P (2) and O (4 x 2 = 8) since it is outside the parenthesis

-The subscript 3 only goes with the Ca

-The coefficient 3 is multiplied to the Ca, P and O after you do the subscripts

-Therefore, this molecule has (3 x 3) Ca+ (3 x 2) P + (3 x 4 x 2) 0

which equals 39 atoms

3Ca3(PO4)2

ionic bonds
Ionic Bonds
  • These are the bonds between a metal and a nonmetal
  • The metal Ion is positively charged and called a cation
  • The nonmetal Ion is negatively charged and called an anion
  • The bonded molecule should be neutrally charged when finished
covalent compounds
Covalent Compounds
  • These can be monatomic or polyatomic compounds
  • It is a bond between two nonmetals
  • They share a pair of electrons
  • They can be subgrouped into polar or nonpolar
  • If a binary compound (2 atoms) – use the same naming rules as in Ionic Compounds
naming covalent compounds
Naming Covalent Compounds

Process:

  • Prefix Indicating # + full name of first

nonmetal

  • Prefix Indicating # + root name of second nonmetal + suffix “ide”
  • Watch for polyatomics and use their proper names
slide26
If it has more than two atoms – need to use the prefixes

Number PrefixNumber Prefix

1 Mono 7Hepta

2 Di 8Octa

3 Tri 9 Nona

4 Tetra 10Deca

5 Penta11Undeca

6 Hexa12Dodeca

for example
For Example:

NOTE THESE ARE ALL NONMETALS WITH NONMETALS!

  • P4S10becomes TetraphosphorousDecasulfide
  • P2O5becomes DiphosphorousPentaoxide
  • SF6 becomes Sulfur Hexafluoride
  • N2O3 is Dinitrogen Trioxide
  • CO is Carbon Monoxide
  • SO2 is Sulfur Dioxide
  • SiBr4 becomes Silicon Tetrabromide
  • Water is really Dihydrogen Monoxide!
slide28
Naming Ionic Compounds is really simple:

1. Name the cation (metal) using its proper name; if it is a polyatomic, do the same

2. Then, using the stem of the anion (nonmetal), simply add the suffix “ide” to it

3. If a transition metal with different possible ions, a roman numeral will tell you which one it is – and it changes the molecular formula!

Examples:

Iron (II) Sulfide = Fe+2 and S-2 combined

Zinc + Chlorine = Zinc Chloride

Iron + Oxygen = Iron Oxide

Lithium + Cyanide = Lithium Cyanide

Ammonium + Fluorine = Ammonium Fluoride

Cobalt + Phosphorous = Cobalt Phosphide

balancing compounds
Balancing Compounds

In an Ionic Compound – balance the molecule using the criss-crossrule. Switch oxidation numbers, making them into subscripts and DROP charges.

Mg +2 + Cl-1

Mg Cl2 The one is understood.

This applies even if using a polyatomic ion

slide30
NH4+ + O-2

(NH4)2O The parentheses are used to keep

the polyatomic together; the 1 is understood

Pb+4+ CO3-2

Pb2 (CO3)4 and this can be simplified by reducing the subscripts to

Pb(CO3)2

chemical equations
Chemical Equations
  • The chemical equation is the rxn formula
  • Reactants  Products
    • Each component will have a phase indicator:
      • (g) meaning it is in its gaseous phase (not just gassy)
      • (l) meaning it is in its liquid phase
      • (s) in its solid phase
      • And (aq) meaning the substance is in a solution of water, aq meaning aqueous
slide32

Must remember which elements are normally diatomic (N2, O2, F2, Cl2, Br2, I2, and H2)

  • All molecules in an equation must be balanced first!!
    • Remember the criss-cross rule!!
  • You may not adjust any subscripts from the original formula
  • You may add and adjust, as you will see, the coefficients in front of each item in the equation
example
Example:
  • H2 + O2 H2O
    • This is the skeleton equation
    • According to the Law of Conservation of Matter, both sides of the arrow must have the SAME number of atoms for each and every element – NO EXCEPTIONS
    • The  can be treated like an = sign
    • In reality, it indicates that some sort of process occurred to cause the reaction
    • So. . . . .
slide34

To balance this simple equation:

    • We ARE NOT ALLOWED TO CHANGE SUBSCRIPTS
    • We CAN ADJUST COEFFICIENTS ONLY
    • The subscripts are the numbers after and below each element’s symbol
    • The coefficients are number in front of a unit (atoms or molecules) and tell how many units there are
    • The coefficients are multiplied out to each and every unit’s atom they are in front of
    • So. . . . .
h 2 o 2 h 2 o
H2 + O2 H2O
  • There are 2 H and 2 O on the reactant side of the equation (the left side)
  • There are 2 H and only 1 O on the product side (the right side)
  • Each side must balance
  • You may add, adjust, finagle, cram, etc. any coefficient in front of any and/or all units to get the equation to balance
  • Therefore: 2H2 + O2 2 H2O
2h 2 o 2 2h 2 o
2H2 + O2 2H2O
  • Now this is balanced!
  • It means it takes 2 hydrogen molecules and one oxygen molecule to form 2 water molecules
balancing equations steps
Balancing Equations Steps:
  • First identify all the reactants and products in the equations
  • Remember – subscripts indicate how many of each element’s atoms are present – with 1 being understood
  • Remember to multiply out all subscripts that are outside a unit in parentheses!
  • YOU CAN’T CHANGE SUBSCRIPTS
  • COEFFICIENTS HAVE TO GO IN FRONT OF A UNIT
slide38

Let’s take the unbalanced equation of:

KClO3 KCl + O2

  • List the elements and how many for both sides of the arrow

K 1  K 1

Cl 1 Cl 1

O 3 O 2

  • Obviously, everything is fine except for oxygen
    • This is where we have to adjust
slide39

We can only use coefficients

    • So we try to multiply each Oxygen by a number to get them to equal out
    • These multipliers become coefficients

K 1  K 1

Cl 1 Cl 1

O 3 x 2 = 6 O 2 x 3 = 6

  • So the new equation is:

2 KClO3  KCl + 3 O2

  • This changes the number of K and Cl now
  • You have to readjust again. . . . . .
2 kclo 3 kcl 3 o 2
2 KClO3 KCl + 3 O2
  • Now we have:

K 2  K 1

Cl 2 Cl 1

O 6 O 6

  • Multiply the product KCl by a coefficient of 2 and it balances
  • Let’s check:

2 KClO3  2KCl + 3O2

K 2  K 2

Cl 2 Cl 2

O 6 O 6

  • It’s Balanced! Finally.
another example
Another Example:

C2H6 + O2 CO2 + H2O

  • List the atoms and numbers:

C 2 C 1

H 6 H 2

O 2 O 2 + 1 = 3

  • Let’s go with C first by multiplying CO2 by a coefficient of 2

C2H6 + O2  2 CO2 + H2O

slide42

This gives us:

C 2 C 1 x 2 = 2

H 6 H 2

O 2 O 4 + 1 = 5

  • Now, let’s balance H by multiplying H2O by 3
  • This gives us: C2H6 + O2 2 CO2 + 3 H2O

C 2 C 2

H 6 H 2 x 3 = 6

O 2 O 4 + 3 = 7

  • It’s still not balanced!
  • Let’s try readjusting Oxygen to get it the same amount
slide43

So, if we change the reactant oxygen to 7 and the product water to 6, we get:

C2H6 + 7 O2 2 CO2 + 6 H2O

  • This also changes our product hydrogen.
  • Therefore, change the reactant C2H6 and the product CO2 to balance and you get:

2 C2H6 + 7O2  4 CO2 + 6 H2O

C 4 C 4

H 12 H 12

O 14 O 14

reaction types
Reaction Types
  • SYNTHESIS (or Direct Combination or Composition) REACTIONS
    • 2 + reactants join together to form a single product
    • Resulting compound is based on common oxidation numbers of the reactant elements
    • There is typically an electron transfer from the element with the lower EN to the one with the higher EN
    • So: A + B  AB or AB + C  ABC
slide45

If two nonmetals involved – a covalent bond formed

  • If two metals – a metallic bond
  • If metal with a nonmetal – ionic bond
slide46

DECOMPOSITION REACTIONS

    • Compounds break down into components
      • AB  A + B or ABC  AB + C
      • Examples. . .

CaCO3 CaO + CO2 Ca(OH)2 CaO + H2O

2 KClO3 2 KCl + 3 O2 2 NaCl 2 Na + Cl2

H2CO3 H2O + CO2

replacement reactions 2 types
REPLACEMENT REACTIONS (2 types)
  • Single Replacement (Displacement Rxn)
    • Key Rule: Metals Replace Metals
      • A + BC  AC + B
    • If Nonmetal – a transfer of e- from more reactive to lesser one
    • Halogens Replace Halogens also
    • Metals replace H in H2O  Metal OH- + H2 (g)
    • Metals replace H in Acids  salt + H2(g)
      • Al + H2SO4  AlSO4 + H2(g)
      • 2 Sc(s) + 6 HCl (aq)  2 ScCl2(aq) + 3 H2(g)
double replacement
DOUBLE REPLACEMENT

Example:

    • FeCl3 + 3 NaOH 3 NaCl + Fe(OH)3

OH goes with FE Cl goes with Na

  • Cations exchange anions with each other
    • No change in oxidation numbers
    • Better know your ions and polyatomics
    • Remember the criss-cross rule and balance each compound after exchanging anions!
    • So: AB + CD  AC + BD
combustion
COMBUSTION
  • An exothermic rxn (gives off energy)
  • Usually find CO2 and H2O in products
  • O2 usually found in reactants
    • CH4(g) + 2 O2 CO2(g) + 2 H2O(g) + heat
    • 2 C4H10(g) + 13 O2(g)  8 CO2(g) + 10 H2O(g)
acid base reactions
ACID/BASE REACTIONS

An acid + base  salt + H2O

  • Acids lose a H+ ion and the bases lose OH- ion
    • These make up one of the products, water
  • Process is called neutralization
  • The produced salt does not have to be NaCl and can be any ionic compound
  • Measure acid with pH scale (1 strong, 7 neutral and 14 is a base)
  • Measure base with pOH scale
acid ph and base poh
Acid (pH) and Base (pOH)
  • Strong Acid (1)

Weak Acid

  • Neutral (7)

Weak Base

  • Strong Base (14)
speed velocity and acceleration
Speed, Velocity and Acceleration
  • The speed of an object is the distance the object travels per unit of time. Speed is a rate which tells you the amount of something that occurs or changes in one unit of time.
  • Speed=distance over time
  • Speed can be divided into two subtitles constant speed & average speed. Constant speed is the speed that does not change. Average speed is the total distance divided by time.
  • Velocity is a speed in a given direction
velocity
Velocity
  • V1 represents the initial or starting velocity
    • If the object starts from a rest, V1 will = 0
  • V2 represents the final velocity of an object
    • If the object ends with a stop, then V2 right at the end will be a zero, but not just a millisecond before that!
    • V = d / t
    • And this means d = vt; and t = d / v
acceleration
Acceleration
  • The acceleration of an object as produced by a net force is directly proportional to the magnitude of the net force, in the same direction as the net force, and inversely proportional to the mass of the object.
  • Acceleration (a) = ΔV / Δt -or-
  • Acceleration = force over mass
newton s 1 st law of motion
Newton's 1st Law of Motion
  • An object at rest tends to stay at rest and an object in motion tends to stay in motion with the same speed and in the same direction unless acted upon by an unbalanced force.
  • Sometimes referred to as the “Law of Inertia."
    • Inertia is the state of rest or resisting a force that may cause motion or a change in velocity
  • Frame of Reference – how the observer sees the change in velocity
  • Frame of Reference (Point of View) can be stationary or moving depending on the observer
  • Example: When a car stops suddenly, all the loose objects will continue forward until they hit something that stops them (have you ever had coffee do this at a stoplight?)
newton s 2 nd law of motion
Newton's 2nd Law of Motion
  • The second law states that the acceleration of an object is dependent upon two variables - the net force acting upon the object and the mass of the object.
  • It explains the relation of force, mass & acceleration.
  • Force=mass x acceleration (F = ma)
  • Weight is also a force = m x g
  • The net force on an object is equal to the product of its acceleration and its mass.
force
Force
  • Force is measured in the SI unit called a Newton (N)
    • 1 N = 1 kg x 1 m / s2

1 N = .225 lbs

1 lb. = 4.448 N

      • Forces usually are in equilibrium (balanced)
      • Weight is a Force (wt = m g)
slide58
Force Continued. . .
  • By definition it is a push or pull
  • It can be divided into two subsets: unbalanced and balanced
  • Unbalanced force can cause an object to start or stop moving; or change its acceleration, velocity or direction
  • A balanced force is equal forces on an object that will not change the object’s motion
acceleration due to gravity
Acceleration - Due to Gravity
  • agrav or just plain g, has a value of 9.80665 m/s2
    • We’ll round this off to 9.81 m/s2
      • Use 10 for guesstimating!
    • Believe it or not – agrav at the equator is 9.7804 m/s2 and at the poles it is 9.8321 m/s2
free fall acceleration
Free Fall Acceleration
  • If v1 (initial velocity is zero or the object is at rest then falls):
    • V2 = gt
    • V2 = √2gh
    • H = ½ gt2
    • H = v2 t

2

if v 1 does not equal zero the object is thrown down or is shot downwards
If v1 does not equal zero. . . the object is thrown down or is shot downwards
  • V2 = v1 + g t
  • V22 = v12 + 2 g h
  • H = v1t + ½ g t2
  • H = v2 + v1 t

2

momentum
Momentum (ρ)
  • Momentum is the product of an object’s mass and velocity
  • It is directly proportional to mass and velocity
  • It’s the tendency for an object to keep in motion
    • p = m v
    • F t = m v; where F t is the impulse or change in momentum
newton s 3 rd law
Newton’s 3rd Law
  • Basically – the law means that for every action there is an equal and opposite reaction
  • A rocket launch – the Fthrustdownwards (action) forces the rocket upwards (reaction) against the Fgravity
  • Remember: Action  Equal/Opposite Reaction
vectors
Vectors
  • Properties of Vectors
    • Vectors can be rearranged into a diagram
    • Size and Direction can not be changed
    • Use the tip to tail method of rearranging vector arrows – addition
    • To subtract vectors – add one to its opposite
    • Example: One Dimension

A(6m) (5 m) + B (6 m) = R (Resultant) of 11 m to the right (same direction = addition)

A (4m) + B (-3m) = R of 1 m to

the left (opposite directions = subtraction)

some key concepts
Some Key Concepts
  • Mass – the amount of matter something has
  • Weight – mass affected by the force of gravity (m x g) – this a Force!
  • Density – how much mass per volume

(d = m / volume)

Can be determined through math or through the displacement of fluid method

*(Remember Archimedes and the crown)

work power energy
Work, Power, Energy

Work

  • Work is a force applied to an object that causes displacement
  • W = F Δ d
    • Measured in Newton-meters (Nm) or Joules (J)
    • Kg m2 / s2 is also called a Joule

Power

  • Power is the rate at which work is being done
  • Measured in Joules per Second or Watts (W) and 1 J/s = 1 W
  • Power (P) = Work / Time
    • P = W / t = F d / t

Energy

  • Potential Energy
    • The stored energy of position, inertia, or ability to do work
    • PE = m g h
  • Kinetic Energy
    • Energy of motion
    • KE = .5 m v2
    • KE = F d
ke pe example
KE / PE Example
  • As KE increases, PE decreases and versa vice
  • As it moves upwards, height increases and PE increases; as it moves downwards, velocity increases and KE increases
  • The pendulum at the two highest points have high PE, but no KE until it starts to move towards the center again.
  • Then, the PE decreases until the bob hits the bottom and KE is at its highest
slide68

No medium needed

  • All are transverse waves
  • Have an electrical and magnetic field at right angles to each other
  • Longest Wavelengths Shortest λ
  • Lowest f Highest f
  • Lowest energy (eV) Highest eV
  • Velocity is the same throughout = c = 300 000 km/s
slide70

Differences between Gravitational and Electromagnetic radiation

There are two principal differences between gravity and electromagnetism, each with its own set of consequences for the nature and information content of its radiation, as described.

  • Gravity is a weak force, but has only one sign of charge.Electromagnetism is much stronger, but comes in two opposing signs of charge.This is the most significant difference between gravity and electromagnetism, and is the main reason why we perceive these two phenomena so differently.
  • Significant Gravitational fields are generated by accumulating bulk concentrations of matter. Electromagnetic fields are generated by slight imbalances caused by small (often microscopic) separations of charge.
  • Gravitational charge is equivalent to inertia.Electromagnetic charge is unrelated to inertia.
slide71

SOUND Equations

  • Vsound = 331.5 + .61 (Co)If no temp given, assume 343 m/s
  • v= d / t
  • d = v t
  • t = d / v

 Denser the material, faster the sound!

  • f = v / λ In Hertz (Hz)
  • λ = v / f
  • v = f λ
  • v = λ / T (period)
  • Intensity (I) = Power (P) / Area (A)
    • Intensity (I) = P / 4 π r2In Watts / meter2
  • Power = I (4 π r2) In Watts
  • Doppler Effect = fo = (v + vd / v + vs) fs
slide76

Magnitude of Charge

  • Coulomb’s Law
  • FElectric = K q q’
  • r2
  • K = 8.988 x 109 Nm2/c2
  • q and q’ are charges of objects
  • r is distance between objects

Coulomb (c) and Amperage (I)

  • Amount of charge flowing through a wire in 1 second with a current of 1 ampere
  • Ampere is 1 Coulomb per second, the intensity (I) of the electrical current
  • Based on the charge of an electron
  • 1 coulomb = 6.242 x 1018 e-
    • Current (I) = Q / t in amperes
      • Measuring the intensity of the electric current
  • Charge of an electron (e-) = 1.60218 x 10-19 c = 1 eV

Potential Difference (V)

  • Amount of work in an electric field to take the charge of 1 coulomb from one point to another
  • Volt is the potential difference across a conductor that carries a current of 1 amp
  • V = W / Q
    • V is potential difference in Volts
      • One volt = J/c
    • W is work done in Joules
    • Q is charge in Coulombs
slide77

Resistance (R)

  • Measured in Ohms  Ω
  • R = V / I
    • I ohm (Ω) = 1 V / Amp
  • Ohm’s Law  V = I R
    • Voltage = Current in Amps x Resistance in Ohms
  • Resistance in Series
    • R1 + R2 + R3 + …. = RTotal
  • Resistance in Parallel
    • 1/R1 + 1/R2 + 1/R3 + …. = 1/RTotal

Capacitance (C)

  • C = Q / V
  • Measured in farads (1 coulomb per volt)
  • Parallel Capacitance
    • C1 + C2 + C3 + … = CTotal
  • Series Capacitance
    • 1/C1 + 1/C2 + 1/C3 + … = 1/CTotal

Work and Power

  • Work (WE) = q V
    • In Joules
  • Power (P) = V q / t
    • In Watts (J/s)
    • Power also = V I = I2 R = V2 / R
magnetism
Magnetism
  • Based on charges of atom’s particles
  • It is a field force – line go from N to S (Faraday Lines)
    • Measured in Teslas or Gauss (1T = 100000G);
    • Earth = .0001T
  • All magnets have two poles – if cut it makes new poles!
  • Can lose magnetism if it is heated past material’s “Curie Temperature” and it returns when cooled
  • Types of Magnetism:
    • Diamagnetic: no magnetism in material
    • Paramagnetic: magnetic only when in a magnetic field
    • Ferromagnetic: due to e- sea model of metal, it can be permanently magnetized