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```Chapter 2 part 1
1
Law of conservation of mass

Dalton

Mass is neither created or destroyed in
chemical rxns.
2
Law of definite proportions
Dalton
 In a given compound the relative
number and kind of atoms is constant.

 Example:
Water = 2 H and 1 O
3
Dalton’s Atomic Theory
1.
2.
3.
4.
each element is composed of atoms
All atoms of a given element are
identical
A chemical compound is the result of
the combination of atoms of 2+
different elements
Atoms of an element are not changed
into different types of atoms by
chemical rxns
6
Millikan
Discovered the charge and mass of an
electron
1.6 X 10 -19 C (coulomb)

7

 Alpha
α
γ
 Gamma
 Beta
β
+2 charge
no charge
-1 charge
8
Rutherford


Most of the mass of
an atom and all of
its positive charge
resides in a very
small dense region
(nucleus)
Most of the volume
of the atom is empty
space around the
nucleus
9
10
Atomic Structure



Proton
 Symbol p+
 Charge = 1.602 x 10-19 C (coulomb)
 Mass = 1.0073 amu
Electron
 Symbol = e Charge = -1.602 x 10-19 C
 Mass = 5.486 x 10-4 amu
Neutron
 Symbol = n0
 no charge
 Mass = 1.0087 amu
11
Units = amu or
g/mol
12
Complete Chemical Symbol
A
X
Z
A = Atomic mass ( p+ n)
Z = Atomic number ( p)
X = element Symbol
13
Question
How many neutrons, protons and electrons are
in carbon?
12
6
C
C = Carbon
12 = Atomic Mass ( protons + neutrons)
6 = atomic Number (protons)
14
12 = p + n
6 = protons
12 – 6 = 6 neutrons
In an atom with out a charge the number of
protons is equal to the number of neutrons
Thus: 6 electrons
15
Question

Write the complete chemical symbol for
the following elements.
 Magnesium
 Sodium
 Tungsten
16
Determining Sub Atomic Particles
12
6
C
C = Carbon
12 = Atomic Mass/weight ( protons + neutrons)
6 = atomic Number (protons)
17
Isotopes

Isotopes
 Atoms
of a given element that differ in number of
neutrons and thus in mass.
Isotope
Protons
Electrons
Neutrons
11C
6
6
5
12C
6
6
6
18
Isotopes Cont.
 When
writing isotopes the
atomic number ( # of p+) stays
the same, but the Atomic Mass
(p+ + n0) changes due to the
neutrons.
19
Formulas

Chemical Indicates actual numbers and
type of atoms in a molecule
 H2O2

Empirical gives only relative number of
atoms of each type
 HO

C2H4
CH2
Structural individual bonds are shown,
indicated by lines
20
average atomic mass
Amu = Average Atomic Mass Unit
 The average atomic mass (weight) of an
element is equal to the sum of the
products of each isotope’s mass (in amu)
multiplied by it’s relative abundance.

21
EXAMPLE OF AVERAGE
ATOMIC MASS PROBLEM

Naturally occurring chlorine is 75.53% Cl35 which has an atomic mass of 34.969
amu, and 24.47% Cl-37, which has an
atomic mass of 36.966 amu.

Calculate the average atomic mass of
chlorine.
22
EXAMPLE OF AVERAGE ATOMIC
MASS PROBLEM (CONT)





Average atomic Mass = [ (%/100) (Atomic Mass) ]
Average atomic mass = (0.7553) (34.969 amu) +
(0.2447) (36.966 amu)
= 26.41 amu + 9.045 amu
= 35.46 amu
NOTE: The average atomic mass of an element
is closest in value to the atomic mass of the
most abundant isotope.
23
Halogens
Alkali
metals
TRANSITION
Alkaline EarthMetalloids
METALS
Metal
Metals
Non Metals
24
DIATOMIC SEVEN
25
Nomenclature
Chapter 2 part 2
Check out these videos for more help or if you are absent
Naming molecular compounds
Writing formulas for molecular compounds
Naming ionic compounds video
Writing formulas for Ionic compounds
26
2.7 Ions and ionic compounds

When negative electrons are removed or
added to an atom the charge of that atom
changes from its neutral state to a charged
state ( + or - )

Ion: charged particle
27
a cation is a particle that carries a positive
electrical charge. The cation gets this
positive charge from losing negatively
charged electrons.
28
Anions are ions that carry a
negative electrical charge. Anions
get their negative charge by
gaining one or more electrons
29
Trick
Na+
Cl-
30
Example
Na = atomic number 11
atomic mass = 23
11 protons, 12 neutrons, 11 electrons
Na- = 11 protons, 12 neutrons, 12 electrons
creating a negatively charged particle.
11 + -12 = -1 charge on Na
31
Reverse example
Na = atomic number 11 atomic mass = 23
11 protons, 12 neutrons, 11 electrons
Na+ = 11 protons, 12 neutrons, 10 electrons
We subtracted a negative charge (e-) creating a
positively charged particle.
11 + -10 = +1 charge on Na
32
Question

How many protons and electrons does
Se2- ion posses?

How many protons and electrons does
Cr3+ ion posses?

What kind of ions are these molecules and
why?
33








Se2- :
anion
34 protons (atomic number)
36 electrons = -34 + -2 = -36
Cr3+ :
Cation
24 protons (atomic number)
21 electrons = 24 – 3 = 21
34
Nomenclature (Naming)

As of 2007 there are 31,000,000 known
compounds.

A: memorize all 31,000,000 names
B: Learn how to name them memorize about 50
things that will allow you to name all 31,000,000.
35
Polyatomic Ions

Atoms joined as a molecule, but they have
a net positive or negative charge.
Example:
NO3-, SO42
I will give you a list. You need to try to
memorize them all… yes all.
36
Polyatomic Ion Rap
37
Cation and Anions to Memorize
+/-3 +/-4
Write these on your binder periodic table
38
Putting the pieces together
Na+
Cl-
NaCl is an ionic
compound
Mg2+
2Cl-
MgCl2 is an ionic
compound
39
Ionic Compounds

Contain both positively and negatively
charged ions.

In general ionic compounds are made of
metals and nonmetals.
40
Covalent compounds

2 negatively charged elements

2 non metals
41
42
43
Question

Write the ionic compound for:

Magnesium and Nitrogen

Magnesium and NO3-
44

Mg3N2

Mg(NO3)2

Count the total number of each type of
atom in the molecules above.
45
Homework

Pg 71

#’s 37,40,45,47,48
46
2.8 Naming ionic compounds

PINK SHEET
47
48
49
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