Mass Spectroscopy of the Elements - AP Chemistry

The denotation is element name-proton+neutron. With 7 protons and 8 neutrons, this necomes Nitrogen-15.

For the carbonate ion, there are 3 resonance structures, and each of the 3 oxygen atoms carries a formal charge of –1 in two out of the three structures. This means that there is a –2/3 charge on each O atom. There are no formal charges in H2O or O2. and in O3, the central O has a positive charge and each outer O has a negative charge, so there is an average charge of –1/2 on the two outer ones.

An element is defined by the number of protons it contains. An isotope is the same element with a different number of neutrons; only two choices have six protons, which carbon has. One of them is 13C, which is the same as the question; the other is 12C.

Note that the question asks for an isotope to 13C. By definition, our answer choice cannot be 13C, but must be a variation (an isotope). We are looking for a second answer with 6 protons, in order to maintain the elemental identity of carbon, but a different number of neutrons (not 7). This makes our final answer a carbon atom with 7 protons, 6 neutrons, or 12C, because it is an isotope to 13C.

isotopes have more neutrons, but their protons are the same

The number following the compound is the isotope number, giving the number of protons and neutrons the element contains per atom. To get the number of neutrons you take the isotope number and subtract the atomic number of that element.

In an atomic representation, the number of neutrons and protons is the upper number
and the number of protons is the lower number. If only one number is given, we can assume
that the number of protons is its atomic number. Therefore (a) 7 protons and 8 neutrons
(b) 47 protons and 62 neutrons (c) 14 protons and 16 neutrons. The total number is 86.

O2- is an ion, and not an isotope; an isotope has the same number of protons and electrons, but a different number of neutrons

0.70 (50 amu) + 0.25 (35 amu) + 0.05 (70 amu) = 47 amu = 47 g/mol

Protons, neutrons, and electrons are all subatomic particles. Alpha particles consist of 2 protons and 2 neutrons bound together in a particle. It is usually emitted during fission reactions.

N has 5 valence electrons. In the ion, N has 4 bonds, and thus 8 bonding electrons and 0 nonbonding electrons. Thus, formal charge = 5- (1/2)*8 - 0 = 1.

Isotopes are examples of the same element existing with different amounts of nuetrons. When an atom gains a neutron, a different isotope of this atom has been formed. Ions are atoms that have either completely lost or completely gained an electron when compared to their ground state configuration. Here the atom has lost an electron and is thus a positive ion.

In this simplistic case, we only have two isotopes, so we know the following:

Where "n" represents the relative abundance of X-59, and "(1-n)" is the relative abundance of X-61. We will algebraically solve for "n":

simplifies to

simplifies to

which leads to

If "n" is equal to 0.515, that means the relative abundance of X-59 is 51.5%, and thus the relative abundance of X-61 is 48.5%.

A calculation would show that the relative abundance of each is 50%, you can see that 8.000 g/mol is exactly the midpoint between the masses of 7 and 9, so each has to be present in equal quantities.

The abundance of will be referred to as "n," and the abundance of is "1-n." Together, we know that these values account for the full abundance of chemical X in existence.

The abundance of is 96.25%, and the abundance of is 3.75%.

Neon has an electron configuration of 1s22s22p6, just like O2-. Sodium has one extra electron, while magnesium has two extra electrons, beyond the neon configuration. Finally, the electron configuration for Cl- is 1s22s22p63s23p6. It is one period below flourine; F- would be isoelectric with neon.

Electrons are notably smaller than nucleons (protons and neutrons), however, they have an equal amount of charge to protons. This is why an atom with the same amount of protons as electrons is considered electrically neutral.

The nucleus itself is positively charged, as the electrons are located relatively far away from the atomic center.

This answer becomes a lot easier if we use some quick mental math. First, notice that if the two isotopes were equal in nature (50/50), than a mole would weigh 130 grams.

Since the number is less than that, we can conclude that isotope X is more dominant than isotope Y, lowering the average mass. Now, we can use math to find the exact percentage.

Start by realizing that isotope X will make up some percentage of the element in nature (we will call this percentage "A") while isotope Y will make up the rest of that percentage (1-A).

This means that 80% of the element is composed of isotope X, and the other 20% is composed of isotope Y. Going back to our mental math, we see that, just as we predicted, isotope X is present in a larger amount than isotope Y.

Isotopes are members of the same element that have a different number of neutrons, and therefore a different atomic weight. The number shown in the top left is always the number of protons plus neutrons (atomic weight), so we can show that an element is an isotope by increasing or decreasing that number.

There are three main isotopes of carbon: carbon-12, carbon-13, and carbon-14. The identity of the element must not change, only the atomic weight. The other answer choices are other types of related chemicals; note that ions are not isomers.

In order to find the percentage of the lighter isotope in the natural sample, we need to use an equation that compares the percentage of each isotope and the natural molar mass. Since both of the isotopes together will account for 100% of the natural sample, we can say that the lighter isotope's percentage is , while the remaining percentage is .

Use this equation to solve for .

In other words, we can conclude that the lighter isotope makes up 40% of the natural sample.