Atomic Structure. Chemistry Mr. McKenzie

Transcription

1 Atomic Structure Chemistry Mr. McKenzie

2 How was the understanding of the atom developed? John Dalton ( ) - developed a model to explain observations made at the time 1. Elements are made of tiny particles called atoms. 2. All atoms of a given element are identical. 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. 5. Atoms are indivisible in chemical processes (not created or destroyed).

3 J.J. Thomson ( ) - Applied an electrical field to a cathode ray Showed that atoms contain negative particles, now called electrons (Noble Prize 1906) Reasoned that there must also be a positive particle to balance out the negative Viewed the atom as uniform distribution of positive and negative particles ( plum pudding model )

4 Ernest Rutherford ( ) - Shot α-particles (+ charged particles) at gold foil and noticed most went straight through and some were deflected Concluded that... The positive charge of an atom was concentrated in a dense center due to the presence of protons The electrons were found outside of the nucleus in mostly empty space

5 James Chadwick ( ) - along with Rutherford, credited with the discovery of the neutron (Noble Prize in 1935) Particle found in the nucleus that contributes to the mass of the atom yet does not have a charge (electrically neutral)

6 Summary of Subatomic Particles Subatomic Particle Location Relative Charge Relative Mass Proton Nucleus + 1 Neutron Nucleus No Charge 1 Electron Electron Cloud - ~1/2000

7 How do atoms differ? Although all atoms are composed of protons, neutrons and electrons, the number of each subatomic particle can vary Each element has its own unique number of protons (the number of protons determines the element) The difference in protons and electrons determines the charge of the atom The number of neutrons in the nucleus impacts the mass of the atom The components of each are represented using atomic symbols +/- charge

8 Write the atomic symbol using the following combinations of subatomic particles: A X #+/ Z 6 protons, 6 neutrons, 6 electrons 15 protons, 16 neutrons, 15 electrons 15 protons, 16 neutrons, 18 electrons Determine the number of protons, neutrons and electrons from the following symbols: 23 Na Ca 18 Cl 27 Al

9 When atoms have a different number of protons and electrons, they take on a charge and are called ions Cations - more protons than electrons; + Anions - more electrons than protons; - The only particle that can change and affect the charge is the number of electrons; if the number of protons changes then the atom is another element

10 When the number of neutrons varies between atoms of the same element, the mass differs and physical properties are slightly altered Isotopes - atoms of the same element that have different numbers of neutrons (protons have to be the same!!!) Some isotopes are unstable and decay/are radioactive (more on this when we cover nuclear chemistry) Carbon-14-8 neutrons (compared to the stable carbon-12 with 6 neutrons); used in determining the age of organic objects Iodine neutrons (compared to the stable iodine-127 with 74 neutrons); used as a radiotracer for thyroid function

11 Where are the electrons located in the atom? We know the electrons are found outside of the nucleus in the electron cloud (mostly empty space due to Rutherford s experiment) Information about the electrons location is based on what happens when atoms gain energy and emit electromagnetic radiation

12 Electromagnetic radiation is a spectrum of various energies with specific wavelengths Wavelength - distance between two consecutive wave peaks Wavelengths have distinct regions that are familiar A short wavelength = high energy (high frequency) A long wavelength = low energy (low frequency)

13 When electrons absorb energy, they move from a ground state (where they started) to an excited state (an area farther away from the nucleus) The excited state is unstable and the electron falls back down to the ground state, emitting the energy equal to what it originally absorbed in the form of a photon of light

14 When the emitted light is analyzed, only certain wavelengths are present, suggesting that the electrons must be moving between distinct regions/levels (b) and not to any region/level (a) If the electrons could go anywhere, all wavelengths of light would be emitted, resulting in a continuous spectrum (all colors) Because the electrons only go to certain areas, only certain wavelengths are emitted Each element produces its own unique line spectrum (only certain colors) when electrons fall

15 Using a Bohr model of hydrogen, we can begin to visualize where the electrons are located and what is happening when they are excited Energy levels (regions where electrons are found) are larger and represent higher energies the farther away from the nucleus Different jumps and falls result in different wavelengths being emitted Farther from the nucleus Farther from the nucleus

16 The Bohr model of the atom (simple circular orbits) is insufficient to explain all elements The wave mechanical model helped explain how electrons are situated in all elements Electrons have a probability of being found in a certain area (they do not circle the nucleus like the Bohr model suggested)

17 Principle Energy Levels and Sub-Levels The primary regions where electrons can be found are the principle energy levels Within each principle energy level are sub-levels Because each energy level gets larger the farther away from the nucleus, each higher principle energy level can have more sub-levels Farther from the nucleus

18 Sub-Levels can come in different types called orbitals An orbital is a region that can hold a maximum of 2 electrons s-orbital - spherical shape; holds a total of two e - ; found in all energy levels

19 p-orbital - dumb-bell in shape; a sub-level consists of 3 different p-orbitals along the x, y and z axis; each holds 2 e- (a total of 6 e- found in the sub-level); found in energy level 2 and higher d-orbital - clover-leaf in shape; a sub-level will hold 5 d-orbitals (a total of 10 e-); found in energy level 3 and higher f-orbital - a sub-level will hold 7 f-orbitals (a total of 14 e-); found in energy level 4 and higher

20 Principle Energy Level # of Sub- Levels Sub-Level (Orbital) Types 1 1 s 2 2 s, p 3 3 s, p, d 4 4 s, p, d, f

21 How can the electron arrangement of atoms be represented? Electrons will always fill the lowest energy orbitals before filling higher energy levels (Aufbau principle); this allows for two ways to represent where the electrons are located in atoms Electron Configurations - shows the distribution of electrons in orbitals; follows an order of orbital filling You do not need to explain why this order happens, just know that it happens!!!

22 Write the electron configuration for the following atoms/ions: Li C Cl Br Br - Ca Ca 2+

23 Orbital diagrams (box diagrams) - show electron configurations using boxes to represent orbitals, grouped by sub-level Each orbital (box) holds a maximum of 2 electrons, which are represented with arrows pointing in opposite directions (to represent the spin of the electron, Pauli exclusion principle) For p, d and f orbitals, each orbital will contain 1 electron before they fill with 2 (lower energies and the repulsion of the electrons)

24 1 box = 1 s orbital arrows point opposite direction to indicate spin 3 boxes = 3 orbitals each orbital holds 1 before they hold 2

25 Write the electron configuration and orbital diagram for the following atoms/ions: C Ar Cl Cl - K + Br Kr