struktur atom kimia kelas x chapter 4 lengkap

struktur atom kimia kelas x chapter 4 lengkap

1. Chapter 4

2. Atom Review
Atoms – the fundamental building blocks of matter.
Major Subatomic Particles
Particle Charge Location Mass
proton +1 nucleus 1 u
neutron 0 nucleus 1u
electron -1 electron
cloud
1/1837 u

3. electrons in shells
nucleus

4. Nucleus – dense central region of atom; contains
essentially all of the atom’s mass.
Atomic Number (Z) – the number of protons in
the nucleus. Each element has a different atomic
number.
The atomic number identifies the element.
Isotopes – atoms of the same element that have
different masses. Isotopes have different masses
because they have a different number of neutrons.

5. Mass Number – the number of protons AND neutrons
in the nucleus.

6. The isotopes of a particular element have
different mass numbers:
Atomic Number: 1 1 1
Mass Number: 1 2 3

7. You can write the mass numbers of isotopes 2 ways:
1) Write the name or symbol of the element followed by
a hyphen and the mass number.
Hydrogen-1 or H-1
Hydrogen-2 or H-2
Carbon-12 or C-12
Carbon-13 or C-13
U-238 Most common isotope of Uranium
(99.3 %)

8. 2) Write the mass number as a superscript
to the left of the chemical symbol:
1
H
2
H
12
C
14
C
Sometimes the atomic number
is added as a subscript:
235
U Fissionable isotope
of Uranium
14
C
6
235
U
92

9. +
N
N
+
-
-
proton
electron
neutron
Shell
Helium-4 isotope
*number of electrons = number of protons
in neutral atom

10. ATOMIC STRUCTUREATOMIC STRUCTURE
HeHe
the number of protons and
neutrons in an atom
44
Mass number
the number of protons in an atom
22
Atomic number
2 electrons (electrons = number of protons
in neutral atom)

11. Ions have a different number of electrons than
the neutral atom.
Negative ions have additional electrons (equal to their charge).
EX: O has 8 electrons; O
2-
has 10 electrons
Positive ions have fewer electrons (equal to their charge).
EX: Al has 13 electrons; Al
3+
has 10 electrons

12. Symbol: Sr P
Charge: 0 0 +2
Atomic #: 20
protons: 16
electrons: 18
neutrons: 50 20
Mass #: 33 31

13. Atomic Mass
Chemists have defined the carbon-12 atom as
having a mass of 12 atomic mass units (u).
Therefore, 1 u = 1/12 the mass of a carbon-12
atom.
1 u is approximately the mass of a single proton
or neutron.

14. Information in the Periodic Table
The number at
the bottom of
each box is
the average
atomic mass
of that
element.
This number is the weighted average mass of all the
naturally occurring isotopes of that element.

15. Average Atomic Mass
1)Almost all elements occur as a mixture of isotopes
2)The percentage of each isotope is a constant
EX: Hydrogen is composed of three isotopes:
Isotope Atomic Mass Natural Abundance
H-1 1.007825 99.985 %
H-2 2.0140 0.015 %
H-3 3.01605 trace

16. The atomic mass listed on the periodic table for
each element is a weighted average of the masses
of the isotopes for that element.
A weighted average takes into consideration the
percentage of each isotope.
Calculate the average atomic mass of hydrogen
using the previous given isotopic data.

17. EX: Chlorine is composed of two isotopes:
Isotope Atomic Mass Natural Abundance
Cl-35 34.9689 75.77 %
Cl-37 36.9659 24.23 %
Calculate the average atomic mass of chlorine.
Calculating Average Atomic Masses

18. ***To calculate average atomic mass:***
• a) Multiply the atomic mass of each isotope by
its percentage
(don’t forget to divide percentage by 100)
• b) Add results of part “a” together to get average
EX2: Copper is composed of two isotopes:
Isotope Atomic Mass Natural Abundance
Cu-63 62.930 69.17 %
Cu-65 64.928 30.83 %
Calculate the average atomic mass of copper.

19. Calculate the average atomic mass of carbon.
EX3: Carbon consists of the following isotopes:
Isotope Atomic Mass Natural Abundance
C-12 12.000 98.89 %
C-13 13.00335 1.11 %
C-14 14.003241 0 (trace amount)

20. Using the information below, calculate the
average atomic mass of element X:

Isotope Percent Abundance Mass (u)

26
X 79.5 25.9

24
X 8.60 24.2

22
X 11.9 22.3