A HISTORY OF THE STRUCTURE OF THE ATOM

A HISTORY OF THE STRUCTURE OF THE ATOM
Unit 2 - Lecture 1: Structure of the Atom A HISTORY OF THE STRUCTURE OF THE ATOM

Let’s begin with… Greek Philosopher Democritus (460-370 B.C.):
all matter composed of small atoms atomos = indivisible What did Democritus conclude about cutting matter in half? There was a limit to how far you could divide matter. You would eventually end up with a piece of matter that could not be cut. Why weren’t Democritus’s ideas accepted? Aristotle was a very famous Greek philosopher who believed that matter could be divided into smaller and smaller pieces forever. He held a very strong influence on popular belief and his views on this were accepted for two thousand years.

1803, John Dalton: atoms are the fundamental building blocks of matter
He performed many experiments to study how elements join together to form new substances He found that they combine in specific ratios and he supposed it was because the elements are made of atoms.

Dalton's Postulates Each element is composed of extremely small particles called atoms. Atoms are indivisible and indestructible particles. Figure 2.1 John Dalton ( )

Dalton's Postulates 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. Figure 2.1 John Dalton ( )

Dalton's Postulates 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. Figure 2.1 John Dalton ( )

Dalton’s Postulates Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

John Dalton’s Atomic Theory (ca 1803)
Unit 2 - Lecture 1: Structure of the Atom John Dalton’s Atomic Theory (ca 1803) Each element is composed of extremely small particles called atoms. All atoms of a given element are identical. The atoms of different elements are different and have different properties (including different masses). Atoms of an element are not changed into different types of atoms by chemical reactions. Atoms are neither created nor destroyed in chemical reactions. This is the Law of Conservation of Mass. Compounds are formed when atoms of more than one element combine. A given compound always has the same relative number and kind of atoms. This is the Law of Definite Composition. 11

John Dalton’s Atomic Theory
Unit 2 - Lecture 1: Structure of the Atom John Dalton’s Atomic Theory Almost right. A good start. very small Structure of the atom after Dalton (ca. 1810)

Unit 2 - Lecture 1: Structure of the Atom
J.J. Thomson (1897): Cathode Rays Atoms subjected to high voltages give off cathode rays. I can demonstrate this for you!

JJ Thomson: https://youtu.be/oddjdB0qfMg What did he discover?
A particle that has MASS and has a NEGATIVE CHARGE The mass was 2000x smaller than the smallest known atom (Hydrogen) This particle must be FROM an atom! Named it the electron

J.J. Thomson: Summarizing what he saw Thomson concluded that the negative charges came from within the atom. A particle smaller than an atom had to exist. The atom was divisible! Since the gas was known to be neutral, having no charge, he reasoned that there must be positively charged particles in the atom. But he could never find them. Electrons are in atoms.

J.J. Thomson – The Electron “Plum pudding” model: Negative electrons are embedded in a positively charged mass. Electrons (-) Unlike electrical charges attract, and that is what holds the atom together. Positively charged mass Structure of the atom after Thomson (ca. 1900)

Radioactivity Radioactivity is the spontaneous emission of radiation by an atom. First observed by Henri Becquerel ( ). Marie and Pierre Curie also studied it. Nobel Prize in 1903 (physics).

Studies of Natural Radioactivity Some atoms naturally emit one or more of the following types of radiation: alpha (α) radiation (later found to be He2+ - helium nucleus) beta (β) radiation (later found to be electrons) gamma (γ) radiation (high energy light) α Alpha particles Electrons (-) γ γ Positively charged mass α Somehow gamma radiation is in there, too. Structure of the atom after Becquerel (early 1900s)

Radioactivity Three types of radiation were discovered by Ernest Rutherford:  particles (positive, charge 2+, mass 7400 times of e-)  particles (negative, charge 1-)  rays (high energy light) Figure 2.8

Robert Millikan Robert Millikan in 1909 determined the size of the charge on an electron. What Millikan did was to put a charge on a tiny drop of oil, and measure how strong an applied electric field had to be in order to stop the oil drop from falling. Since he was able to work out the mass of the oil drop, and he could calculate the force of gravity on one drop, he could then determine the electric charge that the drop must have.

Millkan … By varying the charge on different drops, he noticed that the charge was always a multiple of -1.6 x C, the charge on a single electron. This meant that it was electrons carrying this unit charge. Animation here:

An atomizer sprayed a fine mist of oil droplets into the chamber
An atomizer sprayed a fine mist of oil droplets into the chamber. Some of these tiny droplets fell through a hole in the upper floor. Millikan first let them fall until they reached terminal velocity. Using the microscope, he measured their terminal velocity, and by use of a formula, calculated the mass of each oil drop. Next, Millikan applied a charge to the falling drops by illuminating the bottom chamber with x-rays. This caused the air to become ionized, and electrons to attach themselves to the oil drops. When a drop is suspended, its weight m · g is exactly equal to the electric force applied q · E

Ernest Rutherford (1910) Scattering experiment: Rutherford’s experiment Involved firing a stream of tiny positively charged particles at a thin sheet of gold foil (2000 atoms thick)

The Nuclear Atom Most of the positively charged “bullets” passed right through the gold atoms in the sheet of gold foil without changing course at all. Some of the positively charged “bullets,” however, did bounce away from the gold sheet as if they had hit something solid. He knew that positive charges repel positive charges. (video) Figure 2.11

Rutherford’s ideas This could only mean that the gold atoms in the sheet were mostly open space. Atoms were not a pudding filled with a positively charged material. Rutherford concluded that an atom had a small, dense, positively charged center that repelled his positively charged “bullets.” He called the center of the atom the “nucleus” The nucleus is tiny compared to the atom as a whole.

Ernest Rutherford The Nucleus (and later the Proton) The mass is not spread evenly throughout the atom, but is concentrated in the center, the nucleus. The positively charged part in the atom is the nucleus Electrons (-) are now outside the nucleus. Structure of the atom after Rutherford (1910)

Rutherford’s Model

Let’s talk about Niels Bohr
The Bohr model of the atom Recall: The Rutherford model positive nucleus negative electrons orbiting the nucleus Why don’t the negative electrons collapse

Features of Bohr’s model:
electrons don’t just orbit randomly they are restrict to fixed orbits each orbit represents a different level of energy energy is quantized – i.e., electrons at each energy level possess a specific amount of energy Based his model on the line spectrum of hydrogen

The Line spectrum of Hydrogen
Where do these lines come from??

Hydrogen’s electron absorbs EM energy Jumps to a higher energy level
The electron is in an “excited state” unstable

As the electron falls back down to its normal stable state (“ground state”), it emits the energy it absorbed. Emitted energy is visible as its line spectrum.

Observations Multiple lines but only one electron
The electron can absorb more than one amount of energy Spectrum is not continuous The electron can absorb only specific amounts of energy

Since each energy level is fixed, only certain amounts of energy can be absorbed/emitted. This is why the spectrum emitted is not continuous. Link here:

Bohr’s Model of the Atom
Electrons are “orbiting” around the nucleus. Each level is a specific amount of energy. Electrons can move to higher levels if energy is absorbed by the electron – the electron is temporarily unstable. Unstable electrons fall back down to their “normal” or lower level and release the stored energy as light of a specific wavelength (color).

Bohr’s model had limitations
Was not useful for predicting the line spectra of other elements BUT was a good starting point for modern atomic theory

Discovery of the Neutron

James Chadwick – Neutron (video summary)
Why did we thing there had to be another particle? Something had to be present to help protons from repelling each other. 2 Hydrogen ≠ 1 Helium (1 P + 1 P) (2 P) 4 Hydrogen = 1 Helium Atomic Mass > Atomic number (# P) His Experiment:

James Chadwick – The Neutron In the nucleus with the protons are particles of similar mass but no electrical charge called neutrons. Electrons (-) are now outside the nucleus in quantized energy states called orbitals. (From Niels Bohr and quantum mechanics) The positively charged particles in the nucleus are protons. + n n Structure of the atom after Chadwick (1932)

And a little more to the Atom
Schrodinger & Chadwick (minute 6:50 and minute 8:20)

What we now consider fact about the ATOM…
As provided by MANY experiments and consistent results…

proton (+) neutron electrons - responsible for the volume and size of the atom, negatively charged 10-10 m nucleus - responsible for the mass of the atom, positively charged 10-14 m

If this stadium were the size of an atom’s electron cloud, the nucleus would be the size of a marble setting on the 50 yard line. Electrons occupy the VOLUME, protons and neutrons constitute the MASS of an atom.

Subatomic Particles Protons and electrons are the only particles that have a charge. Protons and neutrons have essentially the same mass. The mass of an electron is so small (2000 time smaller than the proton), we ignore it. Table 2.1 is the way to write this number!

Atomic Facts Feature Size Mass 1 amu = 1 atomic mass unit = x g electron (-) 10-18 m ??? amu + n n proton (+) 10-15 m amu Electrons are outside the nucleus in quantized energy states called orbitals. neutron (0) 10-15 m amu

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: The atomic number (Z)

Atomic Mass The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

Atomic Number Carbon atom The number of protons in the nucleus is called the atomic number Z. Z determines the identity of an element. Saying “the atomic number of an element is 6” is the same as saying “carbon.” The number of electrons in the atom is also Z (because atoms have no net electric charge). How many neutrons are in C? - proton - neutron

Isotopes A 12C Z 6 The number of protons and neutrons in an element is called the mass number A. A = Z + number of neutrons. An element may have different numbers of neutrons but NOT different numbers of protons. Atoms of an element with different numbers of neutrons are called isotopes of that element. - proton - neutron How many neutrons are in C? The answer is “it depends on the isotope.”

Isotopes Isotopes are atoms of the same element with different masses. Isotopes have different numbers of neutrons. 11 6 C 12 6 C 13 6 C 14 6 C

Isotopes number of protons (Z) number of neutrons 6 5 7 8 mass number (A) number of electrons symbol 11 6 11C or C-11 6 12 6 12C or C-12 6 13 6 13C or C-13 6 14 6 14C or C-14 6

Nuclear Symbols Using nuclear symbols to determine the number of p, n, e, and total charge O 16 8 Mass Number = Atomic Number = 16 8 # protons = atomic number = 8 # neutrons = Mass # - Atomic # = = 8 # electrons = # protons = 8

Nuclear Symbols - Atoms
Unit 2 - Lecture 1: Structure of the Atom Nuclear Symbols - Atoms Example: Write the nuclear symbol for the following atoms: 1) 50 p, 70 n 2) 17 e-, 20 n 120Sn 50 37Cl 17

Ions Atoms can gain or lose electrons to become charged particles called ions. A chemical particle that contains a positive or negative charge Cations are positively charged ions. Formed when an atom loses electrons Anions are negatively charged ions. Formed when an atom gains electrons

Ions Formation of a cation 1p e- + Hydrogen atom 1p, 0 n, 1 e- Hydrogen ion (cation) 1p, 0 n, 0 e- 1 H 1 H+ Net charge = 0 Net charge = +1

Ions Formation of an anion 8p 8n 8e- + 2e- 8p 8n 10e- Oxygen atom 8p, 8 n, 8e- Oxygen ion (anion) 8p, 8n, 10e- 16 O2- 16 O Net charge = 0 Net charge = -2

Nuclear Symbols - Ions Practice writing nuclear symbols from information given: 53 p, 74 n, 54 e- 53 proton (= atomic number)  I 74 neutrons + 53 proton  mass number = 127 54 electrons (one more than protons)  1- 127I1- 53

2) 23 e-, 30 n, net charge = +3 # protons
2) 23 e-, 30 n, net charge = +3 # protons? 23 electrons, but charge of 3+ ie 3 more protons than electrons  p= 26  Atomic number = 26  element = Fe 56 3+ Fe 26

Nuclear Symbols Mass Number Charge Atomic Number X Charge = # p - # e-

Nuclear Symbols O 16 8 2- Mass Number = Atomic Number = 16 8 # protons = atomic number = 8 # neutrons = Mass # - Atomic # = = 8 # electrons = # protons - charge = 8 - (-2) = 10

Nuclear Symbols Ba 137 56 2+ Mass Number = Atomic Number = 137 56 # protons = atomic number = 56 # neutrons = Mass # - Atomic # = = 81 # electrons = # protons - charge = 56 - (+2) = 54

Atomic Masses Atomic masses are based on 12C. The mass of 12C (or C-12) is defined to be exactly 12 amu.

Atomic Masses The mass (weight) shown in the periodic table is the mass of the element as its occurs naturally. If the element has more than one isotope, the mass shown is the weighted average of the masses of the isotopes. Mg has 3 isotopes. 24Mg % amu 25Mg % amu 26Mg % amu weighted average of Mg: x 0.1000x 0.1101x 24.31 amu atomic weight of Mg based on natural abundance: amu

If you need more review…
Website - Website - Video -

A HISTORY OF THE STRUCTURE OF THE ATOM

Similar presentations

Presentation on theme: "A HISTORY OF THE STRUCTURE OF THE ATOM"— Presentation transcript:

Similar presentations

About project

Feedback

Log in

Auth with social network:

A HISTORY OF THE STRUCTURE OF THE ATOM

Similar presentations

Presentation on theme: "A HISTORY OF THE STRUCTURE OF THE ATOM"— Presentation transcript:

Similar presentations

About project

Feedback