Showing posts with label atomic structure. Show all posts
Showing posts with label atomic structure. Show all posts
Thursday, October 23, 2008
Tuesday, October 21, 2008
The Review..."Ëlectrons in Atoms"
Hola !!! hello everyone!!! this is Maria scribing for Tuesday's class.... This will be a short recap for the review on the test tomorrow.
Light and Quantized Energy
>Electromagnetic radiation. is a kind of ENERGY that behaves like a(n) WAVE as it travels through space.
--- Light is one type of electromagnetic radiation, others ex. include: X rays,
All waves can be characterized by their wavelength, amplitude, frequency,and speed.
>Wavelength.. is the shortest distance between equivalent point on a continuous wave.
>Amplitude.. is the height of a wave from the origin to a crest or from the origin to a trough.
>Frequency.. is the no. of waves that pass a given point in 1 second.
the photo shows the wavelength and the amplitude.
---A (n) Quantum is the minimum, amount of energy that can be lost or gained by an atom.---
Quantum Theory and the Atom.
1. Ground State. the lowest allowable energy state of an atom.
2. Bohr's Model of the atom predicted the FREQUENCIES of the lines in hydrogen's atomic emission spectrum.
3. According to Bohr's atomic model. >the smaller an electron's orbit, the lower the atom's energy level.
4. Also the larger an electron's orbit, the higher-energy orbit.
5. Bohr proposed that when energy is added to a hydrogen atom, it's ELECTRON moves to a higher- energy orbit.
6. ---- the hydrogen atom emits a photon corresponding to the difference b/w the ENERGY LEVELS associated w/ the 2 orbits transitions b/w.
7. Bohr's atomic model failed to explain the atomic emission spectrum of elements other the Hydrogen.
ex. Bohr's atomic model
QMM (Quantum Mechanical Model ). the modern model of the atom that treats electrons a waves.
Atomic Orbital. a three-dimensional region around around the nucleus representing the probability of finding an electron.
The difference b/w Bohr model and QMM(Quantum Mechanical Model ).
QMM treats electrons as waves and does not describe the electrons paths around the nucleus.
Bohr model treats electrons as particles traveling in specific orbits.
Electron Configurations
>The arrangement of electrons in an atom < style="color: rgb(204, 102, 204);">Ground-State Electron Configuration.
--3 Rules define how electrons can be arranged in an atom's orbital.--
Aufbau Principle. states that each electron occupies the lowest energy orbital available.
Pauli Exclusion Principle. states that a maximum of 2 electrons may occupy a single atomic orbital, but only if the electrons have opposite SPINS.
Hund's Rule. states that a single electrons with the same spin must occupy each equal-energy orbital before additional electrons w/ opposite spins occupy the same orbitals.
Well folks good luck on tomorrows test!!! the next blogger will be...............
Light and Quantized Energy
>Electromagnetic radiation. is a kind of ENERGY that behaves like a(n) WAVE as it travels through space.
--- Light is one type of electromagnetic radiation, others ex. include: X rays,
All waves can be characterized by their wavelength, amplitude, frequency,and speed.
>Wavelength.. is the shortest distance between equivalent point on a continuous wave.
>Amplitude.. is the height of a wave from the origin to a crest or from the origin to a trough.
>Frequency.. is the no. of waves that pass a given point in 1 second.
the photo shows the wavelength and the amplitude.
---A (n) Quantum is the minimum, amount of energy that can be lost or gained by an atom.---
Quantum Theory and the Atom.
1. Ground State. the lowest allowable energy state of an atom.
2. Bohr's Model of the atom predicted the FREQUENCIES of the lines in hydrogen's atomic emission spectrum.
3. According to Bohr's atomic model. >the smaller an electron's orbit, the lower the atom's energy level.
4. Also the larger an electron's orbit, the higher-energy orbit.
5. Bohr proposed that when energy is added to a hydrogen atom, it's ELECTRON moves to a higher- energy orbit.
6. ---- the hydrogen atom emits a photon corresponding to the difference b/w the ENERGY LEVELS associated w/ the 2 orbits transitions b/w.
7. Bohr's atomic model failed to explain the atomic emission spectrum of elements other the Hydrogen.
ex. Bohr's atomic model
QMM (Quantum Mechanical Model ). the modern model of the atom that treats electrons a waves.
Atomic Orbital. a three-dimensional region around around the nucleus representing the probability of finding an electron.
The difference b/w Bohr model and QMM(Quantum Mechanical Model ).
QMM treats electrons as waves and does not describe the electrons paths around the nucleus.
Bohr model treats electrons as particles traveling in specific orbits.
Electron Configurations
>The arrangement of electrons in an atom < style="color: rgb(204, 102, 204);">Ground-State Electron Configuration.
--3 Rules define how electrons can be arranged in an atom's orbital.--
Aufbau Principle. states that each electron occupies the lowest energy orbital available.
Pauli Exclusion Principle. states that a maximum of 2 electrons may occupy a single atomic orbital, but only if the electrons have opposite SPINS.
Hund's Rule. states that a single electrons with the same spin must occupy each equal-energy orbital before additional electrons w/ opposite spins occupy the same orbitals.
Well folks good luck on tomorrows test!!! the next blogger will be...............
Monday, October 20, 2008
Electronegativity worksheets.
Hello Everyone!
In todays wonderful chemistry class we worked on the three worksheets that we had to do over the weekend!
Just in case you missed out on some of the questions, here are the answers :
Worksheet 1 : Determining Electronegativity difference and Percent Ionic Character.
1. a) 2.23 b) 1.27 c) 0.58 d) 1.26 e) 0.68
2. (A) Na and Cl would be considered ionic.
3. % ionic character, electronegativity difference - direct relationship
4. a) 70% b) 36% c) 10% d)33% e)12%
5. ionic percent + covalent percent = 100
6. a) 30% b) 65% c) 90% d) 67% e) 88%
Worksheet 2: Section 6.3 Periodic Trends
1. (C) 2. (C) 3. (D) 4. (B) 5. (A) 6. (A)
7. Ionization energy is the energy required to remove an electron from a gaseous atom.
8. A high ionization energy value is not likely to form a positive ion because it has a strong hold on its electrons.
9. first ionization energies increase as you move from left to right across a period because as you go across the period, the nucleus hold on to its electrons more tightly.
10. In a group, first ionization energies decrease as you move down a group. This decrease in energy happens because atomic size increase as you move down the group. The valence electrons farther from the nucleus, less energy is required to remove them.
11. * don't have.
12. It indicated the relative ability of its atoms to attract electrons in a chemical bond.
13. Electronegativy increase as you move from left to right across a period and decrease as you move down a group.
Worksheet 3: Section 9.5 Electronegativity and Polarity
1. Its the ability to attract electrons in a chemical bond.
2. Fl, 3.98, Halogens, group VII A
3. Fr, 0.7, Alkali metals, group I A
4. Group -> decreases Period -> increases
5. > 1.70 = ionic bond < 70 =" covalent">
6. True
7. False
8. (C) 9. (D) 10. (B) 11. *don't do* 12. (C) 13. (B)
14, 15, 16, 17. *don't do*
Here are the answers to our awesome and fun worksheet! =)
After going over the worksheets we got a review that is needed to be done by tomorrow!
Make sure that everyone studies because the test will be on WEDNESDAY, OCTOBER 22 2008!
GOOD LUCK EVERYONE! =)
next scribe is... M
Sunday, October 19, 2008
Atomic Structure: Electronegativity
Hello everyone, this is Nelsa blogging for Friday's class! This time, I'm going to save after every sentence. I am not going to go through losing everything again.
Soo, let's get to the point.
I. ELECTRONEGATIVITY
electronegativity: Indicates the relative ability of its atoms to attract electrons in a chemical bond.
Basically, as you move across the periodic table - from left to right - electronegativity increases, and as you move down the periodic table, electronegativity decreses.
In grade nine (.. or last year), we were told that a bond between a metal and a non-metal is ionic, and a bond between two non-metals is covalent. That's still true for the most part, but on Friday, we found out why that's true. Chemical bonds between different atoms are never completely ionic or covalent, and what type of bond it is depends on how strongly the bonded atoms attract electrons.
We'll need the electronegativity difference to figure out the character of the bond. Having a difference of 1.70 is considered 50% covalent and 50% ionic. A number that's greater than 1.70 is then considered ionic and a number that's less than 1.70 is then considered covalent.
II. IONIZATION ENERGY
Soo, let's get to the point.
I. ELECTRONEGATIVITY
electronegativity: Indicates the relative ability of its atoms to attract electrons in a chemical bond.
Basically, as you move across the periodic table - from left to right - electronegativity increases, and as you move down the periodic table, electronegativity decreses.
In grade nine (.. or last year), we were told that a bond between a metal and a non-metal is ionic, and a bond between two non-metals is covalent. That's still true for the most part, but on Friday, we found out why that's true. Chemical bonds between different atoms are never completely ionic or covalent, and what type of bond it is depends on how strongly the bonded atoms attract electrons.
We'll need the electronegativity difference to figure out the character of the bond. Having a difference of 1.70 is considered 50% covalent and 50% ionic. A number that's greater than 1.70 is then considered ionic and a number that's less than 1.70 is then considered covalent.
II. IONIZATION ENERGY
ionization energy: The energy required to remove an electron from a gaseous atom.
An atom with a high ionization energy value has a strong hold on its electrons, which means they are less likely to form positive ions. The opposite is true for atoms with a low ionization energy value.
When you move across the periodic table - left to right - the ionization energy increases, and when you move down the table, the energy decreases.
III. ATOMIC RADIUS
The electron cloud surrounding a nucleus doesn't have a clear edge, therefore the atomic radii can't be measure directly.
An atom with a high ionization energy value has a strong hold on its electrons, which means they are less likely to form positive ions. The opposite is true for atoms with a low ionization energy value.
When you move across the periodic table - left to right - the ionization energy increases, and when you move down the table, the energy decreases.
III. ATOMIC RADIUS
The electron cloud surrounding a nucleus doesn't have a clear edge, therefore the atomic radii can't be measure directly.
atomic size: How closely an atom lies to a neighbouring atom.
For metals, the atomic radius is half the distance between adjacent nuclei in an element [d/2], and for elements that occur as molecules, the radius is half the distance between nuclei of identical atoms that are chemically bonded together.
As you move across the periodic table - left to right - the atomic radii decreases, and it increases as you move down a group.
IV. IONIC RADIUS
When atoms lose electrons and form positively charged ions, they become smaller. When atoms gain electrons and form negatively charged ions, they become larger.
Moving across the periodic table - left to right - will decrease the size of positively charged ions, and increase the size of negatively charged ions. The ionic radii of both positive and negative ions increase as you move down the table.
That's basically a summary of what was on those sheets that (you should) have gotten on Friday. Mm, we were given two worksheets for homework, and one that we did in class together.
Yeah.. that's it! I hope you guys had a good weekend. Happy Monday! xD
Next scribe will be ALVINA.
Wednesday, October 15, 2008
Electron Configuration
Okay so its my turn to blog, but today was confusing a bit so i might not do a good job.
Today we learned about electron configuration.
Which is about how energy levels of electrons are designated by principal quantum numbers(n)
each energy level is divided into sublevels. in each level maximum # of electrons is set. we name each level, and the 1st level is called 1s (s for sublevel)
ahhh so hard to explain>>>
in each orbital we can fit 2 electrons.
okay so in "S" we have 1 orbital.
when we reach a higher level we name it differently like 2p which has has 3 orbitals
so the diagram would be.
S= 1 orbital
p= 3 orbitals
d= 5 orbitals
f= 7 orbitals
1s2
1= 1 level
s= number of orbitals
2= number of electrons.
example:
1s2, 2s2
which would mean 1st level has 2 electrons which fills the whole orbital
2nd level had 2 electrons but only fills 1/3 of the orbitals.
I hope this helped cuz i didn't understand it at first.
Angela to scribe tomorrow!
Today we learned about electron configuration.
Which is about how energy levels of electrons are designated by principal quantum numbers(n)
each energy level is divided into sublevels. in each level maximum # of electrons is set. we name each level, and the 1st level is called 1s (s for sublevel)
ahhh so hard to explain>>>
in each orbital we can fit 2 electrons.
okay so in "S" we have 1 orbital.
when we reach a higher level we name it differently like 2p which has has 3 orbitals
so the diagram would be.
S= 1 orbital
p= 3 orbitals
d= 5 orbitals
f= 7 orbitals
1s2
1= 1 level
s= number of orbitals
2= number of electrons.
example:
1s2, 2s2
which would mean 1st level has 2 electrons which fills the whole orbital
2nd level had 2 electrons but only fills 1/3 of the orbitals.
I hope this helped cuz i didn't understand it at first.
Angela to scribe tomorrow!
Atomic Structure
The class did 2 experiments on spectroscopy. They viewed gas discharge tubes through spectroscopes to see the line spectrum that is produced. Hydrogen, helium, mercury, nitrogen, oxygen, and strontium were some elements that were viewed. The line spectra was reproduced on the lab sheet.
Another lab that was done in class was to produce cool blue light. Students mixed luminol, perborate and copper sulfate in a beaker of distilled water. When the copper sulfate was added, the chemical reaction started. With the lights out in the classroom, the colour in the beaker was a blue luminescent one. This was an example of chemiluminscence.
We also covered the historical aspect of the quantum theory-read some bio's of the scientists involved in developing this theory.
Also saw a neat simulation showing an electron absorbing a certain amount of energy (photon) and travelling to an outer orbit. When the electron fell back to a lower orbit it gave off the energy in the form of coloured light.
Another lab that was done in class was to produce cool blue light. Students mixed luminol, perborate and copper sulfate in a beaker of distilled water. When the copper sulfate was added, the chemical reaction started. With the lights out in the classroom, the colour in the beaker was a blue luminescent one. This was an example of chemiluminscence.
We also covered the historical aspect of the quantum theory-read some bio's of the scientists involved in developing this theory.
Also saw a neat simulation showing an electron absorbing a certain amount of energy (photon) and travelling to an outer orbit. When the electron fell back to a lower orbit it gave off the energy in the form of coloured light.
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