Download Chapter 4 Chemical Foundations: Elements, Atoms, and Ions

Survey
yes no Was this document useful for you?
   Thank you for your participation!

* Your assessment is very important for improving the work of artificial intelligence, which forms the content of this project

Document related concepts

History of molecular theory wikipedia, lookup

Unbinilium wikipedia, lookup

Ununennium wikipedia, lookup

Isotopic labeling wikipedia, lookup

Extended periodic table wikipedia, lookup

Tennessine wikipedia, lookup

Periodic table wikipedia, lookup

Livermorium wikipedia, lookup

Chemical element wikipedia, lookup

Dubnium wikipedia, lookup

Oganesson wikipedia, lookup

Isotope wikipedia, lookup

Neptunium wikipedia, lookup

Promethium wikipedia, lookup

Seaborgium wikipedia, lookup

Abundance of the chemical elements wikipedia, lookup

Lawrencium wikipedia, lookup

Einsteinium wikipedia, lookup

Nihonium wikipedia, lookup

Transcript
Dalton’s Atomic Theory
1. Elements are made of tiny particles called atoms.
2. All atoms of a given element are identical (not exactly;
isotopes)
3. The atoms of a given element are different from those of any
other element.
4. Atoms of one element can combine with atoms of other
elements to form compounds. A given compound always
has the same relative numbers and types of atoms.
5. Atoms are indivisible in chemical processes. That is, atoms
are not created or destroyed in chemical reactions. A
chemical reaction simply changes the way the atoms are
grouped together.
1
Laws of Chemical
Combination
Law of Definite Proportions: Different samples of the same
compound always contain its constituent elements in the
same proportions by mass

This is also called the Law of Constant Composition

Suppose we analyze samples of water from different
sources

We will find in each sample the same ratio by mass of hydrogen to
oxygen.
2
So…
 A 10.0 gram sample of water, H2O, is
comprised of 8.9 grams oxygen and 1.1
grams hydrogen.
 Thus, a 100.0 gram sample of water, is
comprised of 89.0 grams of oxygen and 11.0
grams of hydrogen.
3
Laws of Chemical
Combination
Law of Multiple Proportions (or, Dalton’s Law): If two
elements can combine to form more than one compound,
then the masses of one element that combine with a fixed
mass of the other element are in the ratios of small whole
numbers



For instance, the atom carbon forms two stable compounds
with the atom oxygen:
Carbon Monoxide (C1O1) and Carbon Dioxide (C1O2)
Neither C1.2O1.3 nor C1.2O2.2, respectively
4
Formulas of Compounds
Chemical Formula: An expression showing the chemical
composition of a compound in terms of the symbols for the
atoms of the elements involved.
Rules for Writing Formulas
1. Each atom present is represented by its element symbol.
2. Metals are written first, followed by non-metals.
3. The number of each type of atom is indicated by a subscript
written to the right of the element symbol.
4. When only one atom of a given type is present, the subscript
1 is not written.
5
Formulas of Compounds
Write the formula for each of the following compounds:
(a) A molecule contains four phosphorous atoms and ten
oxygen atoms.
(b) A compound contains one uranium atom and six fluorine
atoms.
(c) A compound contains one aluminum atom and three
chlorine atoms.
6
The Structure of the Atom
Atom: The basic unit of an element that can
enter into chemical combination
They are made of even smaller particles called
subatomic particles (Dalton was wrong!)
Electrons
Protons
Neutrons
7
Summary of the
Structure of the Atom
Mass and Charge of Subatomic Particles
Particle
Electron
Mass (grams)
9.1095 x 10-28
Charge Unit
-1
Proton
1.67252 x 10-24
+1
Neutron
1.67495 x 10-24
0
 Nucleus in center of atom – houses protons & neutrons
 Electrons around nucleus in orbitals
8
Atomic Symbols
Atomic Number: The number of protons in the nucleus of an
atom.
Mass Number: The total number of neutrons (variable) and
protons present in the nucleus of an atom.
9
Atomic Symbols
 We use the symbol to represent
the atom
 X = the symbol of the element
 A = the mass number
 Z = the atomic number
 How do we calculate the number
of neutrons?
A
Z
X
 Mass Number =
 Protons + Neutrons
 Thus, neutrons = A - Z
10
Atomic Symbols: Practice
 Give the atomic symbol for a species that
houses 38 protons and 48 neutrons.
 Give the atomic symbol for a species that
houses 38 protons and 50 neutrons.
 What is the difference?
11
Isotopes
 Atoms having the same atomic number but
different mass numbers; i.e., different number
of neutrons.
 The Periodic Table houses average mass
numbers based on percent abundance of each
isotope.
12
Isotopic abundance
 % abundance = (# atoms of isotope/total # of atoms
of all isotopes) x 100%
 For ex: element X has two isotopes
 X-20 and X-21
 Out of every 100 atoms of X, 90 are X-20 and the rest
(10) are X-21
 So, X-20 has an abundance of (90/100) x 100% = 90%
 And X-21 has an abundance of 10%
13
Atomic weight
 Atomic weight = (% abundance
isotope/100)(mass of isotope 1) + …
 An average measurement
 Shown on the periodic table
 Measured in atomic mass units (amu)
 But more commonly in molar mass or
grams/mole
 Will discuss anon
14
Practice:
What’s the atomic weight?
 B-10 = 19.91%
abundance
 10.0129 amu = mass
 B-11 = 80.09%
abundance
 11.0093 amu = mass
19.91%
80.09%
(
)  (10.0129amu)  (
)  (11.0093amu)  10.81amu
100
100
15
Atomic mass unit




Amu or u = 1/12th mass of C-12
= 1.661 x 10-24 g
(C-12= 12 amu)
Used for subatomic particles
16
The mole
 Mole = amt that contains as many “things” as
there are atoms of 12 g of C-12
 1 mole = 6.022 x 1023 particles
 Molar mass (MM) = mass in grams per 1
mole of particles (g/mol)
17
Practice
 How much does one
atom of argon (Ar)
weigh?
1mol
39.948g
-23
(1atom Ar)  (
)

(
)

6.6337x10
g
23
6.022x10 atoms Ar)
1mol
18
Practice
 How many moles and
atoms of Ar are in a
5.00 gram sample?
1mol
5.00g  (
)  0.125mol
39.948g
6.022 x 1023 atoms Ar
0.125mol  (
)  7.53x1022 atoms Ar
1 mol
19
Your turn
 How many grams of niobium (Nb) are there
in a 2.00 mole sample?
 How many atoms?
20
Introduction to the Periodic
Table
21
What do the numbers mean?
 Each cell houses:





Elemental symbol
Elemental name
Atomic number
Molar mass
Sundry information
 Let’s take a look
22
Introduction to the Periodic
Table
 Elements symbolized
 by either one capital letter
 or two letters; first capital, second lower-case






Examples:
B = boron
Cl = chlorine
P = phosphorus
Na = Sodium
Symbols don’t always match up with English names
 Because use Greek and Latin words
23
Introduction to the Periodic
Table
 Groups: The elements in a vertical column of the periodic
table
 “Old” School: Group-A and Group-B
 Taller groups: A
 Shorter groups: B
 “New” School: Groups I - XVIII
 Elements classified in groups based on valence electron
configuration: similar chemical properties
 Group names:





Alkali Metals: The Group 1A elements (LiFr)
Alkaline Earth Metals: The Group 2A elements (BeRa)
Halogens: The Group 7A elements (FAt)
Noble Gases: The Group 8A (HeRn)
Transition Metals: The Group-B elements
24
Introduction to the Periodic
Table
 Period: The elements in each horizontal row
of the periodic table
 There are 7 periods
 Increase in atomic number going to the right
 Two rows at bottom are:
 Lanthanides or rare-earth metals
 Actinides (contain transuranium metals)
25
Introduction to the Periodic
Table
 Elements can be divided into 3 categories:
 Metals
 Metalloids
 Nonmetals




Step-ladder divides metals from nonmetals
Metalloids surround step-ladder
Periodic Table roughly 75% metals
Let’s label the Periodic Table handout!
26
Introduction to the Periodic
Table: Metals
1. Efficient conduction of heat and electricity
2. Malleability (they can be hammered into
thin sheets)
3. Ductility (they can be pulled into wires)
4. A lustrous (shiny) appearance
27
Introduction to the Periodic
Table: Metalloids
 Or semimetals, have properties that fall
between those of metals and nonmetals
28
Introduction to the Periodic
Table: Nonmetals
 Usually poor conductors of heat and electricity
 More varied physical properties than metals
29
Natural States of the
Elements
 All metals and metalloids are solid
 Except mercury = quicksilver
 Non-metals both solid and gaseous
 Noble gases
 H, O, N, F, Cl
 Some nonmetals are diatomic
 A molecule that consists of two similar atoms
 All halogens
 Oxygen, nitrogen, and hydrogen
 One is liquid
 Bromine
30
Natural States of the
Elements
Allotropes: Two or more forms of the same element that differ
significantly in chemical and physical properties.
Example
Phosphorus: red and white
Carbon: Diamond, Graphite, Buckyball
31