Transcript 1- Chapter 2 Notes

Add metal hydride naming
Chapter 2
Atoms, Molecules,
and Ions
Atomic Theory of Matter
The theory that atoms
are the fundamental
building blocks of
matter reemerged in
the early 19th century,
championed by John
Dalton.
Dalton’s Postulates
1. Each element is composed of
extremely small particles called atoms.
Dalton’s Postulates
2. All atoms of a given element are
identical to one another in mass and other
properties, but the atoms of one element
are different from the atoms of all other
elements.
* He was a little wrong on this one…isotopes!
Dalton’s Postulates
3. Atoms of an element are not changed
into atoms of a different element by
chemical reactions; atoms are neither
created nor destroyed in chemical
reactions.
Law of Conservation of Mass
Dalton’s Postulates
4. Compounds are formed when atoms
of more than one element combine; a
given compound always has the same
relative number and kind of atoms.
Law of Definite Proportions
(a.k.a. Law of Constant Composition)
See Chapter 1 Notes
Dalton also deduced the
Law of Multiple Proportions
CO2
1 𝑚𝑜𝑙 𝐶 ∗
CO
12.0 𝑔 𝐶
= 12.0𝑔𝐶
1 𝑚𝑜𝑙𝑒 𝐶
32.0
12.0
16.0𝑔𝑂
2 𝑚𝑜𝑙𝑒 𝑂 ∗
= 32.0𝑔𝑂
1𝑚𝑜𝑙𝑂
1 𝑚𝑜𝑙 𝐶 ∗
= 2.67
12.0 𝑔 𝐶
= 12.0𝑔𝐶
1 𝑚𝑜𝑙𝑒 𝐶
16.0
12.0
= 1.33
16.0𝑔𝑂
1 𝑚𝑜𝑙𝑒 𝑂 ∗
= 16.0𝑔𝑂
1𝑚𝑜𝑙𝑂
2.67
= 2.00 *This is always a whole number!
1.33
The Electron
• J. J. Thompson is credited with their
discovery (1897).
• Streams of negatively charged particles were
found to emanate from cathode tubes.
Millikan Oil Drop Experiment
Determined the
charge &
mass of the
electron
Discovery of the Nucleus
Ernest Rutherford
shot  particles at a
thin sheet of gold foil
and observed the
pattern of scatter of
the particles.
The Nuclear Atom
•
Rutherford
postulated a very
small, dense nucleus
with the electrons
around the outside of
the atom.
• Most of the volume of
the atom is empty
space.
Other Subatomic Particles
• Protons were discovered by Rutherford
in 1919.
• Neutrons were discovered by James
Chadwick in 1932.
Subatomic Particles
(The mass of an electron is so small we ignore it.)
Symbols of Elements
Elements are symbolized by one or two
letters.
Atomic Number
All atoms of the same element have the same
number of protons
Mass Number
The mass number is the total number of
protons and neutrons in the atom.
Isotopes:
• Atoms of the same element with different masses.
• Isotopes have different numbers of neutrons.
11
C
6
12
C
6
13
C
6
14
C
6
Atomic Mass (Weight) of an
Atom
The actual mass can be calculated for any atom:
• Actual Mass (g) = (Mass#)(1.67x10-24g)
• This can be converted to an atomic mass by
relating the mass to carbon-12
Actual Mass (g)
2.00𝑥10−23 𝑔
=
𝐴𝑡𝑜𝑚𝑖𝑐 𝑀𝑎𝑠𝑠 𝑜𝑓 𝑎𝑡𝑜𝑚
12 𝑎𝑚𝑢
Mass of one atom in amu = Mass of a mole of atoms in grams!
*Scientists defined the mole so this would work!
Atomic Mass (Weight) of an
Element
• A weighted average of atomic masses of all the
isotopes of an element.
Use the following data to calculate the atomic
mass for the element Magnesium
Isotope Atomic Mass of Isotope Abundance
Mg - 24
23.982628 µ
78.600 %
Mg - 25
24.963745 µ
10.11 %
Mg - 26
25.960802 µ
11.29 %
(.78600) (23.982628 g) + (.1011) (24.963745 g) +
(.1129) (25.960802 g)
18.850 g
+
2.524 g
24.305 g/mol
+ 2.931 g
Add mass spec here
(see end of ppt)
Periodic Table:
• A systematic
catalog of
elements.
• Elements are
arranged in order
of atomic number.
Periodic Table
• The rows on the
periodic chart are
periods.
• Columns are groups.
• Elements in the same
group have similar
chemical properties.
Groups
These five groups are known by their names.
Periodic Table
Nonmetals
are on the
right side of
the periodic
table (with
the exception
of H).
Periodic Table
Metalloids
border the
stair-step
line (with
the
exception of
Al and Po).
Periodic Table
Metals are
on the left
side of the
chart.
Chemical Formulas
The subscript to the right
of the symbol of an
element tells the number
of atoms of that element
in one molecule of the
compound.
Types of Formulas
• Empirical formulas give the lowest
whole-number ratio of atoms of each
element in a compound.
• Molecular formulas give the exact
number of atoms of each element in a
compound.
• Molecular Formula: C6H12O6
• Empirical Formula: CH2O
Ions
• When atoms lose or gain electrons, they
become ions.
 Cations are positive
 Anions are negative
Ionic Bonds
Ionic compounds (such as NaCl) are
generally formed between metals and
nonmetals.
*See handout on naming and writing formulas!
Molecular Compounds
Molecular compounds
are composed of
molecules and almost
always contain only
nonmetals.
*See handout on naming and writing formulas!
Diatomic Molecules
These seven elements occur naturally as
molecules containing two atoms.
HONClBrIF
Acids
• Compounds containing hydrogen with
something that looks like a negative ion
(for now at least…..)
*See handout on naming and writing formulas!
Organic Molecules (Alkanes)
• Organic molecules contain carbon.
• Alkanes are the most simple organic
compounds. They have chains of
carbon with only single bonds
surrounded by hydrogen.
• They are named by the number of
carbons in the chain + the suffix -ane
Alkanes
# of carbons
Name of alkane
1
Methane
2
Ethane
3
Propane
4
Butane
5
Pentane
6
Hexane
7
Heptane
8
Octane
9
Nonane
10
Decane
Alcohol
• Other classifications of organic
compounds have multiple bonds or
other atoms or groups of atoms
replacing a hydrogen.
• Alcohol has an –OH group instead of a
–H off one of the carbons in an alkane
• They are named with the same root as
the alkanes + the suffix -ol
Starting with propanol they include a
number in front of the name indicating
which carbon the –OH is bonded to
Mass Spectroscopy
Atomic Mass (Weight) of an
Element
• A weighted average of atomic masses of all the
isotopes of an element.
Use the following data to calculate the atomic
mass for the element Magnesium
Isotope Atomic Mass of Isotope Abundance
Mg - 24
23.982628 µ
78.600 %
Mg - 25
24.963745 µ
10.11 %
Mg - 26
25.960802 µ
11.29 %
(.78600) (23.982628 g) + (.1011) (24.963745 g) +
(.1129) (25.960802 g)
18.850 g
+
2.524 g
24.305 g/mol
+ 2.931 g
Where does that data come
from?
• Mass Spectroscopy
Atoms that go into the mass spectrometer
are organized by mass and the relative
amounts are measured.
This is how we found isotopes and
disproved part (the only part so far) of
Dalton’s Atomic Theory.
Mass Spec for Molybdenium
http://www.chemguide.co.uk/analysis/masspec/howitworks.html#top
𝑚𝑎𝑠𝑠
𝑐ℎ𝑎𝑟𝑔𝑒
- we will assume the charge is +1…so this is
the atomic mass
Estimating Atomic Mass of an
Element from its Mass Spectrum
100
23
This scale is usually made by setting the tallest line at 100
All we really care about is this ratio…
if we have 123 atoms, 23 would be B-10 and
100 would be B-11
• You can either find the % abundance
23
(
123
∗ 100% 𝑎𝑛𝑑
100
123
∗ 100%)
and proceed as you learned 1st year.
Or just use the ratios as they are…
23
123
100
10 +
123
11 = 10.9