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Atomic Structure February 2012.

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Presentation on theme: "Atomic Structure February 2012."— Presentation transcript:

1 Atomic Structure February 2012

2 Lesson 1 Explain the law of conservation of mass, the law of definite proportions, and the law of multiple proportions. Summarize the five essential points of Dalton’s atomic theory. Explain the relationship between Dalton’s atomic theory and the law of conservation of mass, the law of definite proportions, and the law of multiple proportions.

3 Write these definitions in your journal.
Law of definite proportions: a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound • Law of multiple proportions: if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers

4 Dalton’s Atomic Theory
All matter is composed of extremely small particles called atoms. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties. Atoms cannot be subdivided, created, or destroyed.

5 Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
In chemical reactions, atoms are combined, separated, or rearranged.

6 Not all aspects of Dalton's theory have proven
to be correct. We now know that: •Atoms are divisible into even smaller particles. •A given element can have atoms with different masses. Some important concepts remain unchanged. • All matter is composed of atoms. • Atoms of any one element differ in properties from atoms of another element.

7 Lesson 2 Summarize the observed properties of cathode rays that led to the discovery of the electron. Summarize the experiment carried out by Rutherford and his co-workers that led to the discovery of the nucleus.

8 Discovery of the Electron
•Experiments in the late 1800s showed that cathode rays were composed of negatively charged particles. •These particles were named electrons.

9 Charge and Mass of the Electron
•Joseph John Thomson’s cathode-ray tube experiments measured the charge-to-mass ratio of an electron. Thomson's Experiment •Robert A. Millikan’s oil drop experiment measured the charge of an electron. Millikan's Experiment •With this information, scientists were able to determine the mass of an electron.

10 Discovery of the Atomic Nucleus
•More detail of the atom’s structure was provided in 1911 by Ernest Rutherford and his associates Hans Geiger and Ernest Marsden. •The results of their gold foil experiment led to the discovery of a very densely packed bundle of matter with a positive electric charge. •Rutherford called this positive bundle of matter the nucleus. Gold Foil Experiment

11 The Structure of the Atom
•An atom is the smallest particle of an element that retains the chemical properties of that element. •The nucleus is a very small region located at the center of an atom. •The nucleus is made up of at least one positively charged particle called a proton and usually one or more neutral particles called neutrons.

12 •Surrounding the nucleus is a region occupied by negatively charged particles called electrons.
•Protons, neutrons, and electrons are often referred to as subatomic particles.

13 Composition of the Atomic Nucleus
Except for the nucleus of the simplest type of hydrogen atom, all atomic nuclei are made of protons and neutrons. A proton has a positive charge equal in magnitude to the negative charge of an electron. Atoms are electrically neutral because they contain equal numbers of protons and electrons. A neutron is electrically neutral.

14 The nuclei of atoms of different elements differ in their number of protons and therefore in the amount of positive charge they possess. Thus, the number of protons determines that atom’s identity.

15 Lesson 3 Discuss the significance of the line-emission spectrum of hydrogen to the development of the atomic model. Describe the Bohr model of the hydrogen atom.

16 The Particle Description of Light
A quantum of energy is the minimum quantity of energy that can be lost or gained by an atom. German physicist Max Planck proposed the following relationship between a quantum of energy and the frequency of radiation: E = hv E is the energy, in joules, of a quantum of radiation, v is the frequency, in s−1, of the radiation emitted, and h is a fundamental physical constant now known as Planck’s constant; h =  1034 J• s.

17 Quantization of energy Energy of a photon
A photon is a particle of electromagnetic radiation having zero mass and carrying a quantum of energy. The energy of a particular photon depends on the frequency of the radiation. Ephoton = hv Quantization of energy Energy of a photon

18 The Hydrogen-Atom Line-Emission Spectrum
The lowest energy state of an atom is its ground state. A state in which an atom has a higher potential energy than it has in its ground state is an excited state.

19 When investigators passed electric current through a vacuum tube containing hydrogen gas at low pressure, they observed the emission of a characteristic pinkish glow. When a narrow beam of the emitted light was shined through a prism, it was separated into four specific colors of the visible spectrum. The four bands of light were part of what is known as hydrogen’s line-emission spectrum.


21 Bohr Model of the Hydrogen Atom
Absorption and Emission Spectra Niels Bohr proposed a hydrogen-atom model that linked the atom’s electron to photon emission. According to the model, the electron can circle the nucleus only in allowed paths, or orbits. The energy of the electron is higher when the electron is in orbits that are successively farther from the nucleus.

22 Bohr Model of the Atom When an electron falls to a lower energy level, a photon is emitted, and the process is called emission. Energy must be added to an atom in order to move an electron from a lower energy level to a higher energy level. This process is called absorption.

23 Photon Emission and Absorption

24 Comparing Atomic Models

25 Lesson 4 Compare and contrast the Bohr model and the quantum model of the atom. Explain how the Heisenberg uncertainty principle and the Schrödinger wave equation led to the idea of atomic orbitals. List the four quantum numbers and describe their significance. Relate the number of sublevels corresponding to each of an atom’s main energy levels, the number of orbitals per sublevel, and the number of orbitals per main energy level.

26 The Heisenberg Uncertainty Principle
German physicist Werner Heisenberg proposed that any attempt to locate a specific electron with a photon knocks the electron off its course. The Heisenberg uncertainty principle states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle. Heisenberg Uncertainty Principle

27 The Schrödinger Wave Equation
In 1926, Austrian physicist Erwin Schrödinger developed an equation that treated electrons in atoms as waves. Together with the Heisenberg uncertainty principle, the Schrödinger wave equation laid the foundation for modern quantum theory. Quantum theory describes mathematically the wave properties of electrons and other very small particles.

28 Electrons do not travel around the nucleus in neat orbits, as Bohr had postulated.
Instead, they exist in certain regions called orbitals. An orbital is a three-dimensional region around the nucleus that indicates the probable location of an electron. Electron cloud

29 Atomic Orbitals and Quantum Numbers
Quantum numbers specify the properties of atomic orbitals and the properties of electrons in orbitals. The principal quantum number, symbolized by n, indicates the main energy level occupied by the electron. The angular momentum quantum number, symbolized by l, indicates the shape of the orbital. Orbitals and Quantum Numbers


31 The magnetic quantum number, symbolized by m, indicates the orientation of an orbital around the nucleus. The spin quantum number has only two possible values—(+1/2 , 1/2)—which indicate the two fundamental spin states of an electron in an orbital.

32 Electrons Accommodated in Energy Levels and Sublevels


34 Quantum Numbers of the First 30 Atomic Orbitals

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