Name: ______Period: ______Date: ______
Notes Chapter 4: The Structure of the Atom
4.1 Early Theories of Matter
A. The Philosophers - believed matter was made of earth, water, air, and fire.
1.Democritus (460-370 BC) was the first to propose that matter was made up of tiny ______he called atomos, which could not be further ______.
a. Without being able to experiment, he could not ______his ideas, and they were ______.
B. John Dalton (1766-1844) – began the development of the modern ______theory.
1. Dalton’s Atomic Theory- 1803
a. All matter is made up of ______.
b. Atoms of ______element are______and are different from those of other elements.
c. Atoms cannot be ______, divided, or ______.
d. Different atoms ______in certain ratios to form ______.
e. In chemical reactions atoms are ______, combined, or ______.
2. Dalton was able to perform ______, observe many ______reactions, and determine the mass ______of elements to verify his theories.
C. Defining the Atom
1. Atom- the smallest particle of an ______that retains the ______of the element.
2. How big is an atom? consider this:
world population in 2000: 6,000,000,000
# of atoms in a single copper penny:
29,000,000,000,000,000,000,000 or ______x ______
3. Nanotechnology – molecular manufacturing which is the atom-by-atom building of ______the size of ______.
4.2 Subatomic Particles and the Nuclear Atom subatomic particles make up ______.
- Discovering the Electron
1. Cathode Ray Tube - Glass ______tube from which ______was removed.a. metal electrodes at each end -______is the negative end, anode is the ______end
2.Cathode Ray- a ray of ______that originates from the cathode and travels to the ______of a cathode ray tube
a. led to the invention of ______and computer monitor images – formed as radiation from the cathode strikes ______-producing chemicals that coat the backside of the ______.
3. By end of the 1800s scientists concluded that
a. No matter what the ______was in the tube or what the electrodes were ______of, the cathode ray was ______which meant that these ______particles were in ______matter.
b. since the particles were attracted to the positive anode, theymust be ______.
c. The ray deflected toward a ______which means it was made of ______ and not just a stream of light.
4. Electrons – ______charged particles.
5. J.J. Thomson(1856-1940) – discovered the first subatomic particle, the ______by determining the mass-to-charge ratio of the particle.
a. Determined that the ______of the charged particle was ______than that of smallest element, ______.
b.Meant that atoms were made of ______particles, disproving part of Dalton’s theory.
c. ______Plum-pudding / Chocolate-chip cookie dough modelof the atom – proposed that negatively charged ______(chips) were distributed through a “dough” of ______charge.
7. Robert Millikan (1868-1953) – 1909 determined that an electron has a ______charge. - 1 (e-)
B. The Nuclear Atom
1. Ernest Rutherford (1871-1937) – discovered existence of ______.
a. used a gold foil experiment to see if positive alpha particles would be ______by the ______in the atom.
b. Since the ______charge was thought to be ______out, he thought it would not alter the path of the alpha particles.
c. Amazingly some were deflected at ______angles which meant there must be a concentrated ______area.
2. Nuclear Model – An atom was mostly ______space through which the electrons move with a tiny ______region called the ______in the center.
3.Nucleus- dense region in ______of atom which is positively charged and contains virtually all of its ______.
4. Neils Bohr (1885 – 1962 ) Electrons ______the nucleus.
a. orbits have a set ______and energy.
b. Lowest energy is the ______orbit.
c. n = ______number or energy level.
d. Radiation (energy) is ______or absorbed when an ______moves from one orbit to another.
e. We see the emitted ______as light or photons.
f. the photons travel at different ______which we see as different ______
5. BUT there was still more ______than could be explained by the protons.
C. Completing the Atom- The Discovery of Protons and Neutrons
1. Rutherford refined the concept of the nucleus to include ______and ______.
2.Proton-subatomic particle carrying a ______charge +1 (p+)
3.James Chadwick (1932) – showed that the nucleaus also contained another subatomic particle.
4. Neutron-has a mass nearly ______to a proton, but carries ______charge (neutral) (n0)
4. Atoms are electrically ______meaning
the number of______= the number of ______.
4.3 How Atoms Differ
~ EachElement is made of ______type of atom. There are 92 natural elements, so there are 92 different kinds of ______.
~ The atoms ______in the number of electrons, protons, and neutrons.
A.Atomic Number- thenumber of ______in an atom.
1. The periodic table is organized left to right and top to bottom by ______atomic number
2.Atoms are neutral and that the number of protons = the number of electrons. So the atomic number will give you the number of ______AND______!!
Atomic number = # protons = # electrons
B. Isotopes and Mass Number
1.While the number of protons must ______the number of electrons, the number of ______may differ.
2.Isotope- atoms with ______number of protons but different numbers of ______.
EX: Potassium (K) has 3 isotopes. All three have ______protons and electrons, but one has ______neutrons, one has _____ neutrons, and one has ______neutrons.
3.Mass Number– is the sum of the number of ______and ______in the nucleus.
Mass Number = # protons + # neutrons
Mass Number – # protons = #neutrons
4.Isotope Identification - the______number is added after an element’s ______to identify the isotope.
EX: Neon-22, Chlorine-35, Uranium-238
5.Symbolic Notation -Shortened type of notation for an element using the chemical ______, atomic number, and the ______number.
The 3 naturally occurring potassium isotopes
Isotopes / Potassium-39 / Potassium-40 / Potassium-41Protons / 19 / 19
Electrons / 19 / 19
Neutrons / 39-19 = 20 / 40-19 = 21 / 41-19 = ______
Symbolic notation / K-41
C. Mass of Individual Atoms
1.Atoms have extremely ______masses which are hard to work with, so scientists use a ______for comparison
2.Standard used is a ______atom
3.Carbon-12 atom has mass of ______atomic mass units
4.atomic mass unit– one (amu) is nearly ______to the mass of 1 proton or 1 neutron.
5.Atomic mass- is the ______average atomic mass of ALL the ______of that element.
D. Calculating Atomic Mass
1. Isotopes of elements exist in nature in ______amounts.
2. Atomic Mass = % ______x atomic mass for each isotope
Then ______all the atomic masses to get the weighted average atomic mass
Ex: Chlorine:
Isotopes
Chlorine-35 exists at 75% so 35 x.75 = ______amu
Chlorine-37 exists at 25% so 37 x .25= ______amu
Weighted Atomic mass is 26.4 + 8.9 = ______amu